GIFT  OF 

W.H.I  vie 


SOLUBILITIES  OF  BASES  AND  SALTS 


K 

Na 

Li 

Ag 

Tl 

Ba 

Sr 

Ca 

Mg 

Zn 

Pb 

Cl 

32.95 
3.9 

35.86 
6.42 

77.79 
13.3 

O.OjlS 
0.059 

0.3 
0.013 

37.24 
1.7 

51.09 
3.0 

73.19 
5.4 

55.81 
6.1 

203.9 
9.2 

1.49 
0.06 

Br 

65.86 
4.6 

88.76 
6.9 

168.7 
12.6 

0.041 
0.066 

0.042 
0.0,15 

103.6 
2.9 

96.52 
3.4 

143.3 
5.2 

103.1 
4.6 

478.2 
9.8 

0.598 
0.02 

I 

137.5 
6.0 

177.9 
8.1 

161.5 
8.5 

0.0.35 
0.071 

0.006 
0.0,17 

201.4 
3.8 

169.2 
3.9 

200 
4.8 

148.2 
4.1 

419 
6.9 

0.08 
0.0,2 

F 

92.56 
12.4 

4.44 
1.06 

0.27 
0.11 

195.4 
13.5 

72.05 
3 

0.16 
0.0,92 

0.012 
0.001 

0.0016 
0.032 

0.0087 
0.0,14 

0.006 
0.0,5 

0.06 
0.002 

NO, 

30.34 
2.6 

83.97 
7.4 

71.43 
7.3 

213.4 
8.4 

8.91 
0.35 

8.74 
0.33 

66.27 
2.7 

121.8 
5.2 

74.31 
4.0 

117.8 
4.7 

51.66 
1.4 

CIO, 

6.6 

0.52 

97.16 
6.4 

313.4 
15.3 

12.25 
0.6 

3.69 
0.13 

35.42 
1.1 

174.9 
4.6 

179.3 
5.3 

126.4 
4.7 

183.9 
6.3 

150.6 
3.16 

BrO, 

6.38 
0.38 

36.67 
2.2 

1525 
8.20 

0.59 
0.025 

0.30 
0.009 

0.8 
0.02 

30.0 
0.9 

85.17 
2.3 

42.86 
1.5 

58.43 
1.8 

1.3 
0.03 

*O, 

7.62 
0.35 

8.33 
0.4 

80.43 
3.84 

0.004 
0.0,14 

0.059 
0.0,16 

0.05 
0.001 

0.25 
0.0,57 

0.25 
0.007 

6.87 
0.26 

0.83 
0.02 

0.002 
0.043 

OH 

SO4 

142.9 
18 

116.4 
21. 

12.04 
5.0 

0.01 
0.001 

40.04 
1.76 

3.7 
0.22 

0.77 
0.063 

0.17 
0.02 

0.001 
0.0,2 

0.0*5 
0.045 

0.01 
0.0,4 

11.11 
0.62 

16.83 
1.15 

35.64 
2.8 

0.55 
0.020 

4.74 
0.09 

0.0323 
0.0410 

0.011 
0.0,6 

020 
0.015 

35.43 
2.8 

53.12 
3.1 

0.0041 
0.0,13 

0,0. 

63.1 
2.7 

61.21 
3.30 

111.6 
6.5 

0.0025 
0.0315 

0.006 
0.0,1 

0.0,35 
0.0414 

0.12 
0.006 

0.4 
0.03 

73.0 
4.3 

.'  .*  ! 

0.042 
0.0B5 

C,04 

30.27 
1.6 

3.34 
0.24 

7.22 
0.69 

0.0034 
0.0S17 

1.48 
0.030 

0.0085 
0.0338 

0.0046 
0.0326 

0.0,55 
0.0443 

0.03 
0.0027 

0.0364 
0.044 

0.0S16 
0.0,54 

CO, 

108.0 
5.9 

19.39 
1.8 

1.3 

0.17 

0.003 
C.0,,1 

4.95 
0.10 

0.0023 
0.0311 

0.0011 
0.047 

0.0013 
0.0313 

0.1 
0.01 

0.004? 
0.0,3? 

0.0,1 
0.043 

The  upper  number  in  each  square  gives  the  number  of  grams  of  the 
anhydrous  salt  held  in  solution  by  100  c.c.  of  water.  The  loiver  number  it 
the  molar  solubility,  i.e.,  the  number  of  moles  contained  in  one  liter  of  the 
saturated  solution.  The  numbers  for  small  solubilities  have  been  abbreviated. 
Thus  0.084  =  0.0000004.  For  some  other  solubilities,  see  page  131. 


****     •  ""-^         . 


GENEEAL  CHEMISTEY 
FOE  COLLEGES 


BY 

ALEXANDER   SMITH 
it 

PROFESSOR  OF  CHEMISTRY,  AND  HEAD  OF  THE  DEPARTMENT, 
COLUMBIA  UNIVERSITY 


SECOND  EDITION 

ENTIRELY    REWRITTEN 


NEW  YORK 

THE  CENTUEY  CO, 

1916 


.    .      .     , 

«• 

:-    , 


\t\\\* 


COPYRIGHT,  1905,  1906,  1908,  1916, 

BY 
THE  CENTURY  CO. 


First  Edition,  May,  1908 

Reprinted  October,  1908;  April,  1909  ;  May,  1910;  May,  1911; 
April,  1912  ;  April,  1913  ;  May,  1914  ;  January,  1915 

Second  Edition,  January,  1916 

Reprinted  July,  1916;  August,  1916; 

September,  1916  ;  October,  1916 


GIFT  OF 


PREFACE   TO   THE   FIRST   EDITION 


THE  present  work  differs  from  the  Author 's  "Introduction  to 
General  Inorganic  Chemistry"  in  being  intended  for  pupils  who 
can  devote  less  time  to  the  study  of  the  science,  and  whose  needs 
can  be  satisfied  by  a  less  extensive  course.  It  resembles  the 
larger  work  in  the  arrangement  of  the  contents  and  in  the  general 
method  of  treatment.  The  matter,  and  particularly  the  theoreti- 
cal matter,  however,  has  been  simplified  and  has  been  confined 
strictly  to  the  most  fundamental  topics.  Such  parts  of  the  theory 
as  are  thus  given,  are  presented  with  the  same  fullness  as  before, 
and  are  illustrated  and  applied  with  all  the  persistence  needed  to 
insure  full  apprehension  and,  ultimately,  spontaneous  employment 
by  the  student.  Such  parts  as  could  not  be  treated  in  this  way, 
within  the  limits  set  by  the  plan  of  the  book,  have  been  omitted. 
Methods  materially  different  from  those  used  in  the  "  Introduction" 
have  been  employed  in  presenting  many  topics.  Conspicuous 
differences  of  this  kind  will  be  noted  particularly  in  the  treatment 
of  combining  proportions,  formulae  and  equations,  molecular  and 
atomic  weights,  chemical  equilibrium,  ionic  substances  and  their 
interactions,  and  the  theory  of  precipitation. 

The  writer  desires  to  express  his  profound  gratitude  to  the  many 
chemists  who  have  made  valuable  criticisms  and  suggestions/ 
Most  of  these  comments  applied  to  the  "Introduction  to  General 
Inorganic  Chemistry,"  but  many  of  them  have  been  used  in 
preparing  this  work  (General  Chemistry  for  Colleges),  and  all  will 
be  considered  in  the  second  edition  of  the  larger  book. 

For  critical  reading  of  the  whole  of  the  proofs  of  the  present 
work,  the  writer  desires  especially  to  thank  Messrs.  A.  T.  McLeod 
and  Alan  W.  C.  Menzies  of  the  University  of  Chicago.  Other  cor- 
rections and  suggestions  will  be  gladly  received  by  the  author. 

ALEXANDER  SMITH. 
Chicago,  April,  1908. 

985148 


PREFACE   TO   THE   SECOND   EDITION 


IN  preparing  the  second  edition,  the  entire  book  has  been  re- 
written. The  introduction  to  the  subject  has  been  improved  and 
greatly  simplified,  and  several  difficult  topics  have  been  trans- 
ferred to  later  chapters.  The  explanations  of  the  theoretical 
subjects,  and  of  the  methods  of  making  calculations  have  been 
clarified  and  additional  illustrations  have  been  given.  In  view 
of  its  importance  to  prospective  students  of  biology  and  medicine, 
osmotic  pressure  is  treated  in  greater  detail.  To  add  greater 
interest  to  the  study  of  the  science,  and  because  of  their  edu- 
cational value,  the  historical  references  have  been  expanded, 
many  more  applications  of  chemistry  have  been  discussed,  and 
the  number  of  figures  has  been  considerably  increased.  Exten- 
sive new  sections  on  oxidation  and  reduction,  and  on  various 
methods  of  writing  equations,  on  radio-activity,  and  on  electro- 
motive chemistry  have  been  added.  Briefer  new  sections  on 
atomic  numbers,  colloids,  foods,  explosives,  water  treatment,  and 
many  other  subjects,  have  been  included.  New  pedagogical  de- 
vices have  been  introduced.  So  far  as  recent  advances  can  be 
apprehended  and  applied  by  first-year  college  students,  the  treat- 
ment has  been  brought  up  to  date. 

The  author  is  greatly  indebted  to  Messrs.  P.  C.  Haeseler,  and 
Herbert  E.  East  lack  for  much  assistance  in  reading  the  proof  and 
for  many  valuable  suggestions. 

ALEXANDER  SMITH. 

New  York,  January,  1916. 


Vll 


CONTENTS 


CHAPTER  PAGE 

I.  THE  CHEMICAL  VIEW  OF  MATTER 1 

II.   CHEMICAL  CHANGE  AND  THE  METHODS  OF  STUDYING  IT.  . .  11 

III.  OXYGEN 25 

IV.  ATOMIC  WEIGHTS,  SYMBOLS,  FORMULAE,  AND  EQUATIONS  ...  40 
V.  HYDROGEN 49 

VI.  VALENCE.     CALCULATIONS 61 

VII.  THE  MEASUREMENT  OF  QUANTITY  IN  GASES.     RELATIONS 

BETWEEN  STRUCTURE  AND  BEHAVIOR  OF  MATTER 70 

VIII.  WATER 85 

IX.   MOLECULAR  WEIGHTS  AND  ATOMIC  WEIGHTS 100 

X.   SOLUTION 121 

XI.   HYDROGEN  CHLORIDE.     CALCULATIONS 141 

XII.   CHLORINE 154 

XIII.   ENERGY  AND  CHEMICAL  CHANGE 167 

XIY-»  CHEMICAL  EQUILIBRIUM 177 

XV.   THE  HALOGEN  FAMILY 192 

XVLx  DISSOCIATION  IN  SOLUTION 210 

XVII.   OZONE  AND  HYDROGEN  PEROXIDE 219 

XVIII.   IONIZATION 226 

XIX.   IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS 245 

XX.   SULPHUR  AND  HYDROGEN  SULPHIDE 264 

XXI.   THE  OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR 275 

XXII.   SELENIUM  AND  TELLURIUM.     THE  CLASSIFICATION  OF  THE 

ELEMENTS • 293 

XXIII.  OXIDES  AND  OXYGEN  ACIDS  OF  THE  HALOGENS.    OXIDA- 

TION AND  REDUCTION 306 

XXIV.  THE  ATMOSPHERE.     THE  HELIUM  FAMILY 328 

XXV.   NITROGEN  AND  AMMONIA 338 

XXVI.   OXIDES  AND  OXYGEN  ACIDS  OF  NITROGEN 347 

XXVII.   PHOSPHORUS 362 

ix 


X  CONTENTS 

CHAPTER  PAGE 

XXVIII.  CARBON  AND  THE  OXIDES  OF  CARBON 375 

XXIX.  THE  HYDROCARBONS.     FLAME 389 

XXX.  THE  CARBOHYDRATES  AND  RELATED  SUBSTANCES 402 

XXXI.  ORGANIC  ACIDS  AND  SALTS.     ALCOHOLS,  ESTERS.    FOODS..  412 

XXXII.  SILICON  AND  BORON 425 

XXXIII.  THE  BASE-FORMING  ELEMENTS 434 

XXXIV.  THE  METALLIC  ELEMENTS  OF  THE  ALKALIES:    POTASSIUM 

AND  AMMONIUM 443 

XXXV.  SODIUM  AND  LITHIUM.     IONIC  EQUILIBRIUM  CONSIDERED 

QUANTITATIVELY 457 

XXXVI.   THE  METALLIC  ELEMENTS  OF  THE  ALKALINE  EARTHS  .  .  .  473 

XXXVII.  COPPER,  SILVER,  GOLD 500 

XXXVIII.  GLUCINUM,  MAGNESIUM,  ZINC,  CADMIUM,  MERCURY.    THE 

RECOGNITION  OF  CATIONS  IN  QUALITATIVE  ANALYSIS  .  .  .  523 

XXXIX.   ELECTROMOTIVE  CHEMISTRY 539 

XL.  ALUMINIUM  AND  THE  METALS  OF  THE  EARTHS 553 

XLI.   GERMANIUM,  TIN,  LEAD 567 

XLII.  ARSENIC,  ANTIMONY,  BISMUTH 582 

XLIII.  THE  CHROMIUM  FAMILY.    RADIUM 595 

XLIV.   MANGANESE 617 

XLV.  IRON,  COBALT,  NICKEL 625 

XLVI.  THE  PLATINUM  METALS 644 

APPENDIX.  648 


GENERAL    CHEMISTRY    FOR   COLLEGES 


GENERAL  CHEMISTRY  FOR  COLLEGES 


CHAPTER  I 
THE  CHEMICAL  VIEW  OP  MATTER 

CHEMISTRY  is  a  science  which  deals  with  all  forms  of  matter.  It 
considers  the  natural  kinds,  such  as  rocks  and  minerals,  as  well  as 
materials  like  fat  and  flour  obtained  from  animals  or  plants.  It 
deals  also  with  artificial  products  like  paints  or  explosives.  When 
we  wish  information  about  any  specimen  or  kind  of  matter,  we 
consult  a  chemist.  Now  chemists  have  worked  out  a  point  of 
view  which  enables  them  to  attack  any  problem  connected  with 
matter  in  a  systematic  manner  and  to  state  the  results  in  a  clear 
and  simple  way.  To  learn  something  of  chemistry,  we  must 
acquire  this  point  of  view  and  master  the  technical  language  the 
chemist  uses  in  stating  and  discussing  his  results. 

Properties.  —  Suppose  that  a  piece  of  rusty  iron  is  submitted 
to  the  chemist.  He  at  once  examines  the  rust  and  notes  that  it 
is  solid,  reddish-brown  in  color  and  earthy  in  appearance.  He 
separates  some  of  it  from  the  iron  and  finds  it  to  be  brittle,  that  is, 
easily  broken  and  capable  of  being  pulverized  in  a  mortar.  He 
finds  that  its  density  is  about  4.5,  that  is  to  say,  1  c.c.  (Appendix  I) 
of  it  weighs  about  4.5  g.  On  heating  some  of  it  in  a  flame,  he  finds 
that  it  does  not  melt,  and  must  therefore  have  a  very  high  melting- 
point.  These  qualities  he  calls  properties,  and  more  especially 
physical  properties.  Since  all  specimens  of  iron-rust  show  exactly 
the  same  properties,  he  often  calls  them  specific  physical  proper- 
ties, because  they  are  properties  shown  by  all  specimens  of  a  par- 
ticular species  of  matter. 

After  removing  any  rust  by  filing  or  scraping,  the  chemist  ex- 
amines the  iron,  and  finds  a  fresh,  clean  surface  to  be  almost  white 
and  metallic  in  appearance.  The  metal  is  tenacious,  so  that  it  can 

1 


COLLEGE    CHEMISTRY 


be  bent  but  not  easily  broken.  It  is  ductile  and  can  therefore  be 
drawn  out;  into  wire.'  He  fir»ds  that  its  density  is  about  7.5,  and 
that  the  metal  is  incapable  of  being  melted  in  an  ordinary  flame. 
In  addition,  he  finds  it  to  be  strongly  attracted  by  a  magnet, 
while  rust  is  not  attracted. 

The  chemist,  then,  studies  what  he  calls  the  specific  physical 
properties  of  each  material,  in  order  that  he  may  be  able  to  recog- 
nize various  materials. 

Substances.  —  All  specimens  of  iron  show  one  set  of  proper- 
ties and  all  specimens  of  iron-rust  show  a  different  set,  peculiar  to 
rust.  The  chemist  calls  any  definite  variety  of  matter,  all  speci- 
mens of  which  show  the  same  properties,  a  substance.  Iron  is  one 
substance  and  rust  another.  A  substance  is  recognized  by  its 
properties. 

The  point  of  view  of  the  chemist  thus  consists  in  describing  any 
material  by  ascertaining  whether  it  is  made  up  of  one,  or  of  more 
than  one  substance.  He  describes  it  by  naming  the  substances 
which,  by  a  study  of  their  properties,  he  has  found  in  it. 

Two  Illustrations  of  the  Study  and  Description  of  Ma- 
terials. —  If  a  piece  of  granite  is  examined  by  a  chemist,  he 
observes  at  once  that  it  is  spotted  in  appearance,  and  made  up  of 
several  crystalline  materials  of  differing  nature.  He  therefore 
breaks  it  up  and  studies  the  properties  of  the  fragments.  Some  of 
the  fragments  of  granite  are  dark  and  with  a  penknife  can  easily 


FIG.  1. 


FIG.  2. 


FIG.  3. 


be  split  into  transparent  sheets,  thinner  than  paper.  These  par- 
ticular fragments  are  in  all  respects  like  mica  (Fig.  1).  This  sub- 
stance is  a  mineral  which,  in  certain  neighborhoods,  occurs  in 
large  masses,  and  sheets  of  it  ("isinglass")  are  used  to  fill  the 
openings  in  stoves.  Others  of  the  fragments  are  clear  like  glass, 
and  are  very  hard  (see  Appendix  II),  and  have  all  the  properties 


THE    CHEMICAL   VIEW    OF   MATTER 


of  quartz  or  rock  crystal  (Fig.  2),  which  is  another  substance 
well  known  to  the  chemist.  The  remaining  fragments  are  less 
clear  than  is  quartz,  and  are  not  so  hard.  They  can  be  split  into 
layers,  but  not  nearly  so  easily  as  can  mica.  They  form  oblong 
crystals,  differing  in  this  also  from  quartz,  which  shows  hexagonal 
crystals.*  This  substance  is  felspar  (Fig.  3).  Thus  the  chemist 
studies  the  physical  properties  of  the  fragments,  and  finds  that 
there  are  three  different  substances  in  granite.  He  reports  that 
the  components  of  granite  are  mica,  quartz,  and  felspar. 

When  flour  is  examined  by  the  chemist,  it  appears  to  the  eye  to 
be  all  alike.  Under  the  microscope,  even,  all  he  can  learn  is  that 
it  consists  largely  of  grains,  which  have  the 
characteristic  appearance  (first  property)  of 
grains  of  starch  (see  Fig.  108,  p.  403).  He 
places  some  flour  on  a  square  piece  of  cheese- 
cloth and  encloses  it  by  tying  with  a  thread 
(Fig.  4).  On  kneading  the  little  bag  in  a  vessel 
of  water,  the  water  becomes  milky.  When  the 
milky  water  stands,  the  white  material  settles 
to  the  bottom,  the  water  can  be  poured  off, 
and  the  deposit  can  be  dried.  This  white  sub-  f 

stance,  when  boiled  with  water,  gives  an  almost  clear  liquid  which 
jellies  on  cooling.  This  is  another  property  of  starch.  A  little 
tincture  of  iodine  (solution  of  iodine  in  alcohol),  dropped  on  a 
part  of  the  starch,  causes  the  latter  to  turn  blue.  This  is  a  very 
characteristic  property  of  (and  therefore  test  for)  starch.  When 
the  bag  of  flour  is  kneaded  persistently  in  water  which  is  frequently 
changed,  the  material  finally  ceases  to  render  the  water  milky. 
The  starch  has  all  been  washed  out.  When  the  bag  is  now  opened, 
a  sticky  material  is  found  in  it.  This  is  called  gluten.  The  chemist 
therefore  finds  that  the  flour  contains  starch  and  gluten.  He 
learns  this  by  separating  the  components. 

Law  of  Component  Substances.  —  Every  material  can  be 
described  as  being  composed  of  one  substance,  or  as  being  a  mixture 

*  Crystals  (see  also  Index)  are  natural  forms,  of  geometrical  outline,  which 
solid  substances  assume.  Usually  each  substance  has  a  more  or  less  distinct 
form  of  its  own,  the  particular  angles  at  which  the  faces  meet  being  peculiar  to 
the  substance.  Its  individual  crystauine  form  is  therefore  a  specific  physical 
property  of  each  substance. 


4  COLLEGE    CHEMISTRY 

of  two  or  more  component  substances,  each  of  which  has  a  definite 
set  of  specific  physical  properties.  This  is  the  first  and  most 
fundamental  law  of  chemistry.  This  conception  was  first  clearly 
stated  by  Lomonossov  (1742),  a  Russian  author,  statesman,  and 
chemist  (1711-1765). 

Mixtures  and  Impurities.  —  A  material  containing  more 
than  one  component  substance  is  called  a  mixture.  The  charac- 
teristic of  a  mixture  is  that  each  of  the  component  substances, 
although  mixed  with  the  others,  possesses  exactly  the  same  prop- 
erties as  if  it  were  present  alone.  No  one  of  the  components 
affects  any  other  component,  or  alters  any  of  its  properties. 
Granite  and  flour  are  typical  mixtures. 

When  a  specimen  is  composed  mainly  of  one  substance,  and 
contains  only  minute  amounts  of  one  or  more  other  substances, 
it  is  frequently  spoken  of  as  a  specimen  of  the  main  substance 
containing  certain  specified  substances  as  impurities.  To  be  called 
an  impurity,  the  foreign  matter  need  not  be  dirty  or  offensive. 
Thus,  common  salt  usually  contains  a  little  magnesium  chloride, 
a  white  crystalline  solid,  as  an  impurity,  and  it  is  this  impurity 
which  becomes  damp  in  wet  weather.  Again,  compounds  of 
lime  and  magnesium  are  common  impurities  in  drinking  water. 

Components.  —  The  ingredients  of  a  mixture  are  called  the 
components  (Lat.,  put  with),  because  they  are  simply  placed  to- 
gether, without  change,  and  can  be  separated  without  change. 

Bodies  or  Specimens.  —  It  will  be  seen  that  substance  is  a 
general  term,  like  the  word  "dog,"  covering  the  whole  species. 
The  substance  iron  includes  all  the  iron  in  the  universe.  When 
we  refer  to  a  particular  piece  of  iron,  we  call  it  a  body  or  a  speci- 
men. If  the  body  is  homogeneous  (all  parts  alike),  it  may  be 
made  of  a  single  substance.  If  it  is  heterogeneous  (differing  in 
different  parts)  it  is  a  specimen  of  a  mixture  like  granite. 

The  Rusting  of  Metals.  —  If  we  return  once  more  to  the 
subject  of  rusty  iron,  we  find  another  point  which  interests  the 
chemist.  If  the  iron  is  kept  moist  —  for  example,  by  lying  in 
the  grass  or  partly  immersed  in  water  —  the  layer  of  rust  gradu- 


THE    CHEMICAL   VIEW   OF   MATTER     .  5 

ally  becomes  thicker,  and  the  core  of  iron  becomes  thinner,  until 
it  finally  disappears.  The  rust  seems  to  be  formed  from  the  iron, 
in  presence  of  air  and  moisture.  The  iron,  particle  by  particle, 
loses  the  properties  of  iron  and  simultaneously  acquires  those  of 
rust.  Now  the  chemist  is  concerned,  not  only  with  recognizing 
substances,  but  also  with  the  ways  in  which  substances  change  and 
new  substances  are  produced. 

Several  other  familiar  metals  rust,  as  does  iron,  but  the  change 
is  slower.  Thus,  lead  rusts  (tarnishes)  slowly,  and  zinc  still  more 
slowly.  The  change  can  be  hastened  by 
heating.  For  example,  if  some  lead  is 
melted  in  a  porcelain  crucible  (Fig.  5) 
and  is  stirred  with  an  iron  wire,  a  dirty 
yellow  powder  collects  on  the  surface. 
Gradually  more  and  more  of  the  powder 
is  formed  and  less  and  less  of  the  metallic 
lead  remains,  until  at  last  all  the  metal  is 
gone.  Melted  tin,  when  treated  in  the 
same  way,  gives  a  white  powder. 

Explanation  of  Rusting.  —  The  first 
fact  which  seemed  to  throw  light  on  the  sub- 
ject was  discovered  by  a  French  physician,  Fia-  5- 
Jean  Rey  (1630),  who  found  that  the  rusts  of  tin  and  lead,  made 
by  heating  and  stirring,  were  heavier  than  the  original  pieces  of 
metal.  He  inferred,  "correctly,  that  the  additional  material  which 
caused  the  increase  in  weight  came  from  the  air.  He  imagined, 
however,  that  the  rust  was  not  a  new  substance,  but  a  sort  of 
froth,  and  therefore  a  mixture  of  air  with  the  metal.  Other  in- 
vestigators, such  as  Hooke  (1635-1703)  and  particularly  Mayow 
(1645-1679),  in  England,  explained  the  increase  in  weight  by  sup- 
posing that  some  material  from  the  air  had  combined  with  the 
metal.  In  other  words,  iron,  for  example,  was  .one  substance 
composed  of  iron  only,  and  rust  was  another  substance,  made  by 
union  of  iron  and  a  material  from  the  air,  and  not  a  mere  mixture. 

It  was  Lomonossov  (1756)  who  first  proved  by  an  experiment 
that  the  extra  material  did  come  from  the  air.  He  placed  some 
tin  in  a  flask,  sealed  up  the  mouth  of  the  vessel,  and  weighed  the 
whole.  The  flask  was  then  heated  and  the  tin  was  converted  into 


6  COLLEGE    CHEMISTRY 

the  white  powder.  So  long  as  the  flask  remained  sealed,  no  change 
in  weight  was  found  to  have  occurred.  When  the  mouth  of  the 
flask  was  opened,  however,  some  air  rushed  in,  and  the  total 
weight  was  then  found  to  be  greater.  Evidently,  during  the  heat- 
ing, a  portion  of  the  original  air  had  forsaken  the  gaseous  condi- 
tion and  joined  itself  to  the  tin  to  form  the  powder.  This  left  a 
partial  vacuum  in  the  flask,  and  more  air  entered  when  the  latter 
was  opened.  Eighteen  years  later  the  same  experiment  was  made 
by  Lavoisier,  who  drew  the  same  conclusion.  The  rusting  of 
other  metals  was  found  to  be  due  to  the  same  cause.  Lavoisier 
named  the  gas,  taken  from  the  air,  oxygen. 

The  conclusion  can  be  confirmed  in  various  ways.  For  example, 
when  the  air  is  pumped  out  of  the  flask  before  it  is  sealed,  the  metal 
can  be  heated  in  the  vacuum  indefinitely  without  rusting. 

Experiment  to  show  the  Nature  of  Rusting.  —  That  a 
part  of  the  air  is  consumed  when  iron  rusts  is  easily  proved.  We 
moisten  the  interior  of  a  test-tube  and  sprinkle  some  powdered 
iron  so  that  it  covers  and  adheres  to  the 
whole  interior  surface.  We  then  set  the 
tube  mouth  downwards  in  a  dish  of  water 
(Fig.  6) .  At  first,  the  pressure  of  the  water 
compresses  the  air  in  the  tube  very  slightly, 
and  the  water  ascends  above  the  mouth  to 
the  extent  of  a  small  fraction  of  an  inch 
only.  As  the  moist  iron  slowly  rusts,  how- 
ever, the  oxygen  is  gradually  removed,  and 
the  pressure  of  the  atmosphere  outside  slowly 

pushes  the  water  farther  up  the  tube.  After  an  hour  or  more,  the 
water  has  ascended  about  one-fifth  of  the  total  distance  towards 
the  top  of  the  tube.  Evidently  part  of  the  air  has  forsaken  the 
gaseous  condition,  and  the  water  has  been  forced  up  to  take  its 
place.  Inspection  now  shows  some  reddish  particles,  where  rusting 
has  taken  place.  The  rust,  then,  is  made  up  of  a  part  of  the  iron 
and  all  of  the  oxygen  that  the  tube  contained. 

Of  course,  much  of  the  iron  powder  is  still  gray,  and  has  not 
rusted.  The  air  in  the  tube  did  not  contain  oxygen  enough  to 
combine  with  all  the  iron.  The  iron  that  remains  is  as  little  able 
to  rust  in  the  remaining  gas  as  in  a  vacuum. 


THE   CHEMICAL   VIEW   OF  MATTER  7 

Incidentally  we  learn  from  this  experiment  that  atmospheric 
air  contains  about  one-fifth  (20  per  cent)  oxygen  by  volume. 
The  remaining  four-fifths  is  almost  all  nitrogen  (79  per  cent),  a 
substance  which  combines  with  very  few  materials,  while  the 
balance  (1  per  cent)  is  made  up  of  gases  which  do  not  enter  into 
combination  with  any  known  substance.  If  lead,  tin,  or  zinc 
had  been  heated  in  an  enclosed  volume  of  air,  they  likewise  would 
have  taken  out  the  20  per  cent  of  oxygen  and  would  have  left 
the  other  gases. 

The  Law  of  Chemical  Change.  —  The  three  examples  of 
rusting  show  that  specimens  of  matter  can  lose  their  original 
properties  and  acquire  new  ones.  Since  a  substance  is  "a  spe- 
cies of  matter,  with  a  constant  set  of  properties,"  we  are  compelled 
to  decide  that,  when  a  material  changes  its  properties,  it  has,  in 
doing  so,  become  a  new  substance.  This  consideration  calls  to 
our  attention  the  second  of  the  fundamental  laws  of  chemistry, 
namely,  that  the  material  forming  one  or  more  substances  (such  as 
oxygen  and  iron),  without  ceasing  to  exist,  may  be  changed  into 
one  or  more  entirely  different  substances.  Such  a  change  is  called 
a  chemical  change,  or  action,  or  interaction,  or  reaction. 

The  commoner  kinds  of  chemical  actions  can  be  divided,  for 
convenience,  into  four  varieties.  We  can  now  define  the  first  of 
these. 

First  Variety  of  Chemical  Change:    Combination.  —  In 

each  case  of  rusting,  two  substances  (a  gas  and  a  metal)  came 
together  to  form  a  third  substance  (an  earthy  powder).  Appar- 
ently two  substances  may  come  together  in  two  different  ways. 
They  may  form  a  mixture,  in  which  both  substances  are  present 
and  retain  their  properties,  or  they  may  come  together  to  form  a 
single  substance  with  different  properties.  When  two  (or  more) 
substances  unite  to  form  one  substance,  the  change  is  called  chem- 
ical combination  or  union.  The  product  is  called  a  compound 
substance. 

We  are  very  careful  never  to  speak  of  a  compound  substance 
as  a  mixture.  Rust  is  not  a  mixture  of  iron  and  oxygen;  it  shows 
none  of  the  properties  of  either.  Nor  do  we  call  a  mixture  (like 
granite)  a  compound,  or  the  operation  of  mixing,  combination  or 


8  COLLEGE    CHEMISTRY 

union.  These  are  technical  words,  in  chemistry  and,  to  avoid 
confusion,  may  be  used  only  with  due  regard  to  their  technical 
meanings. 

Constituents.  —  As  we  have  seen,  we  speak  of  the  substances 
in  a  mixture  as  the  components.  When  we  wish  to  refer  to  the 
forms  of  matter  which  are  chemically  united  in  a  compound,  we 
call  them  the  constituents  (Lat.,  standing  together)  of  the  compound 
substance.  Thus,  iron  and  oxygen  are  the  constituents  of  rust. 

The  chemist  separates  (p.  3)  the  components  of  a  mixture,  for 
that  is  all  that  is  necessary.  He  liberates  the  constituents  of  a 
compound,  however,  because  they  are  bound  together  in  chemical 
combination. 

The  names  given  to  compounds  are  usually  devised  so  as  to 
indicate  the  nature  of  the  constituents.  Thus,  iron-rust  is  oxide 
of  iron  (or  ferric  oxide,  from  Lat.  ferrum,  iron).  The  yellowish 
powder  from  lead  is  lead  oxide  or  oxide  of  lead,  and  the  white 
powder  from  tin  is  oxide  of  tin. 

A  Condensed  Form  of  Statement.  —  We  may  represent  a 
chemical  combination,  or  indeed  any  kind  of  chemical  change,  in  a 
condensed  form,  thus: 

Iron  +  Oxygen  — >  Oxide  of  iron  (ferric  oxide). 

Each  name  stands  for  a  substance.  Two  substances  in  contact 
with  one  another  (mixed),  but  not  united  chemically,  are  con- 
nected by  the  +  sign.  The  arrow  shows  where  the  chemical 
change  comes  in,  and  the  direction  of  the  change.  We  read  the 
statement  thus:  Iron  and  oxygen  brought  together  under  suitable 
conditions  undergo  chemical  change  into  oxide  of  iron,  called  also 
ferric  oxide.  Similarly  we  may  write: 

Lead  +  Oxygen  — >  Oxide  of  lead. 
Tin  +  Oxygen  — >  Oxide  of  tin. 

The  Increase  in  Weight  in  Rusting.  —  As  we  have  seen, 
the  process  of  rusting  is  accompanied  by  a  slow  increase  in  the 
weight  of  the  solid,  due  to  the  gradual  addition  of  oxygen  to  the 
metal.  Now,  this  increase  in  weight  ceases  of  its  own  accord, 
when  a  certain  maximum  has  been  reached.  This  occurs  when 


THE    CHEMICAL  VIEW   OF  MATTER  9 

the  last  particles  of  the  metal  have  disappeared.  Thus,  the  lead 
gains  in  weight  until  every  100  parts  of  the  metal  have  gained  7.72 
parts  of  oxygen,  and  the  tin  until  every  100  parts  have  gained 
26.9  parts  of  oxygen.  When  these  increases  have  occurred,  the 
metal  is  found  to  have  been  all  used  up,  and  prolonged  heating 
and  stirring  cause  no  further  union  with  oxygen  and  no  further 
change  in  weight.  This  fact,  that  each  substance  limits  itself 
of  its  own  accord  to  combining  with  a  fixed  proportion  of  the  other 
substance,  in  forming  a  given  compound,  is  one  of  the  most  strik- 
ing facts  about  chemical  combination.  In  mixtures,  any  propor- 
tions chosen  by  the  experimenter  may  be  used.  In  chemical 
union,  the  experimenter  has  no  choice;  the  proportions  are  de- 
termined by  the  substances  themselves.  Thus,  100  parts  of  iron 
when  turning  into  ordinary  red  rust  take  up  43  parts  of  oxygen, 
no  more  and  no  less. 

This  fact  enables  us  to  make  our  condensed  statements  more 
specific  and  complete  by  including  in  them  the  proportions  by 
weight  used  in  the  chemical  change: 

Iron  (100)    +  Oxygen  (43)     —>  Ferric  oxide  (143). 
Lead  (100)  +  Oxygen  (7.72)  -» Oxide  of  lead  (107.72). 

The  following  numbers,  which  represent  the  same  proportions 
by  weight,  are  the  ones  commonly  used  by  chemists: 

Iron  (111.68)  +  Oxygen  (48)  ->  Ferric  oxide  (159.68). 

Summary.  —  Thus  far,  we  have  learned  that  chemistry  deals 
with  substances  and  their  physical  properties,  and  with  the  changes 
which  substances  undergo.  We  have  discussed  and  defined  a 
number  of  important  words  expressing  fundamental  chemical 
ideas.  Finally,  we  have  touched  upon  the  weights  of  the  ma- 
terials used  in  chemical  change,  a  subject  of  great  importance 
which  will  be  more  fully  developed  in  a  later  chapter. 

Exercises.  —  1.  Take  one  by  one  the  words  or  phrases  printed 
in  black  type  and  the  titles  of  the  sections  in  this  chapter,  and 
endeavor  to  recollect  what  you  have  read  about  each.  In  each 
case  try,  (a)  to  recall  the  meaning  and  to  state  it  in  your  own 
words;  (6)  to  recall  the  facts  associated  with,  and  the  reasoning 
which  lead  up  to  the  point  in  question;  (c)  to  recall  examples 


10  COLLEGE   CHEMISTRY 

illustrating  the  conception  and  to  apply  the  conception  in  detail 
to  each  example.  Whenever  memory  fails  to  give  a  perfectly 
clear  report  of  the  matter  in  hand,  the  text  must  be  read  and 
re-read  until  the  essential  point  can  be  repeated  from  memory. 

Use  the  same  method  in  all  future  chapters.  A  useful  prac- 
tice is  to  employ  a  pencil  as  you  read  and  to  underline  systemati- 
cally all  the  important  facts  and  statements,  and  then  to  go  back 
and  apply  to  each  marked  place  the  process  described  above. 

2.  Define  the  following  terms:   Specific  gravity,  tenacity,  melt- 
ing-point, specific  physical  property,  pure  body,  vacuum. 

3.  Is  it  logical  to  say  "pure  substance?" 

4.  Why  do  we  decide  that  granite  is  a  mixture  and  iron  a 
single  substance? 

5.  Do  the  statements  in  the  text  indicate  that  air  is  a  mixture 
or  a  compound? 

6.  What  weight  of  oxygen  would  be  required  to  convert  25 
grams  of  lead  into  oxide  of  lead? 

7.  Make  a  list  of  the  technical  words  we  have  defined,  and  place 
the  definition  opposite  to  each. 

8. .  What  weight  of  tin  would  be  contained  in  15  grams  of  oxide 
of  tin? 

9.  If  any  of  the  following  are  mixtures,  mention  the  facts  which 
show  them  to  contain  more  than  one  substance:  (a)  muddy  water, 
(6)  an  egg,  (c)  milk. 

10.  In  recognizing  a  specimen  to  be  quartz,  does  the  chemist 
consider  (a)  the  weight,  (b)  the  temperature,  (c)  the  length  of  the 
specimen?     If  not,  why  not? 

11.  Give  a  list  of  the  specific  properties  mentioned  in  this 
chapter. 


CHAPTER  II 

CHEMICAL*  CHANGE  AND  THE  METHODS   OF 
STUDYING  IT 

WE  must  now  take  up  two  new  examples  of  chemical  change. 
They  will  aid  us  in  introducing  one  or  two  additional  conceptions 
and  laws.  These  are  continually  used  by  the  chemist,  and  without 
them  we  cannot  begin  the  systematic  study  of  the  science. 

Another    Case   of   Combination:     Iron   and  Sulphur. — 

Since  oxygen  is  an  invisible  gas,  there  is  a  slight  difficulty  in  real- 
izing that  rusting  consists  in  the  union  of  two  substances  —  this 
gas  and  a  metal.  The  present  example  is  less  interesting  histori- 
cally, but  it  is  simpler  because  both  substances  are  visible  and  are 
easily  handled.  The  case  of  iron  and  sulphur  will  enable  us  to 
illustrate  the  same  point  of  view  and  to  practice  the  application 
of  the  same  technical  words.  It  will  also  introduce  us  to  two 
manipulations  —  nitration  and  evaporation  —  which  are  fre- 
quently used  by  the  chemist. 

We  begin  by  observing  the  physical  properties  of  the  two  sub- 
stances. Those  of  iron  have  already  been  noted  (pp.  1-2).*  Sul- 
phur is  a  pale-yellow  substance  of  low  specific  gravity  (sp.  gr.  2). 
It  is  easily  melted  (m.-p.  112.8°  C.).  It  does  not  dissolve  in  water 
—  that  is,  it  does  not  mix  completely  with. and  disappear  in  water, 
as  sugar  does  on  stirring.  It  does  dissolve  readily  in  carbon 
disulphide,  however.  It  crystallizes  in  rhombic  forms  (Fig.  7). 
It  is  not  attracted  by  a  magnet. 

*  References  to  previous  pages  are  used  in  order  to  save  needless  repetition 
in  writing.  The  beginner  requires  endless  repetition  in  his  reading,  however, 
and  must  form  the  habit  of  examining,  in  conjunction  with  the  current  text,  the 
parts  referred  to.  The  passages  cited  are,  by  the  reference,  made  part  of  the 
current  text,  which  will  usually  not  be  clear  without  them.  The  same  remark 
applies  to  topics  referred  to  by  name.  Such  topics  must  be  sought  in  the  index. 

All  terms,  and  especially  those  borrowed  from  physics,  if  not  perfectly 
familiar,  must  be  looked  up  in  a  work  on  physics  or  in  a  dictionary. 

11 


12 


COLLEGE    CHEMISTRY 


Study   of  the   Mixture,   before    Combination.  —  Now,   if 
some  iron  filings  and  pulverized  sulphur  are  stirred  together  in  a 
mortar,  the  result  is  a  mixture.     True,  the  color  is  not  that  of 
either  substance,  but  with  a  lens  particles 
of  both  substances  can  be  seen.     Passing  a 
magnet  over  the  mixture  will  easily  remove 
a  part  of  the  iron,  and  with  the  help  of  a  lens 
and  a  needle  the  mixture  can  be  picked  apart 
particle  by  particle,  completely.   We  can  sep- 
arate the  components  of  the  mixture  more  ex- 
FlG-  7-  peditiously,  however,  by  using  manipulations 

based  upon  certain  suitable  properties.  Thus,  sulphur  dissolves  in 
carbon  disulphide  while  iron  does  not.  If,  there- 
fore, a  part  of  the  mixture  is  placed  in  a  dry  test- 
tube  along  with  some  carbon  disulphide  (Fig.  8), 
and  is  shaken,  the  liquid  dissolves  the  sulphur 
and  leaves  the  iron.  To  complete  the  separa- 
tion, the  iron  must  be  removed  from  the  liquid 
by  filtration,  and  the  sulphur  recovered  by  evap- 
oration of  the  carbon  disulphide. 

\ggr 

Filtration.  —  Iron,  or  any  solid,  when  it  is 
mixed  with  a  liquid  or  with  a  solution  (like  the  solution  of  sulphur 

in  carbon  disulphide)  is  said  to  be  sus- 
pended in  the  liquid.  If  the  solid  is  one 
that  settles  rapidly,  the  liquid  may  be 
separated  from  the  solid,  in  a  rough 
way,  by  pouring  off  as  much  of  the  clear, 
supernatant  liquid  as  possible.  This  is 
called  decantation. 

A  complete  separation  is  effected  by 
pouring  the  mixture  on.  to  a  cone  of 
filter  paper  supported  in  a  glass  funnel 
(Fig.  9).  The  liquid,  together  with 
anything  that  may  be  dissolved  in  it, 
FlG-  9-  runs  through  the  pores  of  the  paper  and 

down  the  hollow  stem  of  the  funnel.  The  liquid  is  then  called  the 
nitrate.  The  particles  of  the  suspended  solid  are  too  large  to  pass 
through  the  pores,  and  so  collect  on  the  surface  of  the  filter  paper. 


CHEMICAL   CHANGE   AND   METHODS   OF   STUDYING   IT      13 

This  operation,  like  everything  the  chemist  does,  takes  advantage 
of  the  physical  properties  of  the  various  materials. 

The  material  remaining  on  the  paper  (the  residue),  when  dry, 
is  wholly  attracted  by  a  magnet  and  shows  all  the  other  properties 
of  iron. 

Evaporation.  —  To  recover  the  sulphur,  the  solution  in  carbon 
disulphide  —  the  nitrate  —  is  poured  into  a  porcelain  evaporat- 
ing dish  (Inflammable!  Keep  flames  away).  When  the  vessel 
is  set  aside,  the  liquid  gradually  passes  off  in  vapor  (e-vapor-ates) . 
Sulphur,  however,  does  not  evaporate  at  room  temperature  and 
remains  as  a  residue,  in  the  form 
of  crystals  of  rhombic  outline  in 
the  bottom  of  the  dish  (Fig.  10). 
Here,  again,  physical  properties 
have  been  utilized. 

Since  the  physical  properties 

of  two  substances  are  not  changed  by  mixing,  we  have  thus  used 
the  properties  of  the  iron  and  sulphur  so  as  to  separate  them  once 
more.  The  iron  is  on  the  paper;  the  sulphur  is  in  the  dish. 

Combination  of  Iron  and  Sulphur.  —  But  iron  and  sulphur 
are  capable  of  combining.  If  we  alter  the  conditions  by  raising 
the  temperature  of  some  of  the  dry  mixture,  as  we  did  in  causing 
lead  to  rust  rapidly,  chemical  union  sets  in.  When  we  place  some 
of  the  original  mixture  of  iron  and  sulphur  into  a  clean  test-tube 
and  warm  it,  we  soon  notice  a  rather  violent  development  of  heat 
taking  place,  the  contents  begin  to  glow,  and  what  appears  to  be  a 
form  of  combustion  spreads  through  the  mass.  The  heating  em- 
ployed at  the  start  falls  far  short  of  accounting  for  the  much 
greater  heat  produced.  When  these  phenomena  have  ceased, 
and  the  test-tube  has  been  allowed  to  cool,  we  find  that  it  now 
contains  a  somewhat  porous-looking,  black  solid.  This  material 
is  brittle;  it  is  not  magnetic;  it  does  not  dissolve  in  carbon  di- 
sulphide; and  close  examination,  even  under  a  microscope,  does 
not  reveal  the  presence  of  different  kinds  of  matter.  This  sub- 
stance is  known  to  chemists  as  ferrous  sulphide  and,  as  we  see,  its 
properties  are  entirely  different  from  those  of  its  constituents. 

In  this  connection  we  must  not  omit  to  notice  that,  as  in  rusting, 


14  COLLEGE   CHEMISTRY 

a  certain  fixed  proportion  will  be  used  in  forming  the  compound. 
We  find  that,  for  7  parts  of  iron,  almost  exactly  4  parts  by  weight 
of  sulphur  are  required.  If  more  iron  is  put  into  the  original 
mixture,  then  £ome  unused  iron  will  be  found  in  the  mass  after  the 
action.  If  too  much  sulphur  is  employed,  some  may  be  driven  off 
as  vapor  by  the  heat  and  any  that  remains,  beyond  the  correct 
proportion,  can  be  dissolved  out  of  the  ferrous  sulphide  with  car- 
bon disulphide.  The  sulphur  which  has  combined  with  the  iron, 
however,  is  no  longer  present  as  sulphur  —  it  has  no  longer  the 
properties  of  sulphur,  and  therefore  cannot  be  dissolved  out: 
Iron  (55.84)  +  Sulphur  (32.07)  ->  Ferrous  sulphide  (87.91). 

Another  Illustration:  Mercuric  Oxide.  —  It  has  long  been 
known  that  air  contains  an  active  and  an  inactive  gas.  The 
Chinese  called  them  yin  and  yang,  respectively.  Mayow  (1643- 
1679)  showed  that  the  active  gas  caused  rusting,  that  it  was 
absorbed  by  paint  (really  by  the  linseed  oil)  in  "drying,"  that  it 
supported  combustion  of  wood  and  sulphur,  and  that  it  is  neces- 
sary to  life,  being  absorbed  by  the  blood  from  the  air  entering  the 
lungs.  It  was  not  until  1774,  however,  that  a  pure  specimen  of 
this  gas  was  obtained,  by  Priestley,  and  was  rec- 
ognized to  be  a  special  kind  of  gas  different  from 
ordinary  air.  The  gas  (later  to  be  named  oxygen) 
was  made  by  Priestley  from  mercuric  oxide,  a 
bright  red,  rather  heavy  powder.  When  the  oxide 
is  heated  (Fig.  11),  we  find  that  a  gas  is  given  off. 
This  gas  is  easily  shown  to  be  different  from  air, 
since  a  glowing  splinter  of  wood  is  instantly  re- 
FIO.  11.  lighted  on  being  immersed  in  it.  The  gas  is  pure 

oxygen.  During  the  heating,  we  notice  also  that  a  metallic  coating 
appears  on  the  sides  of  the  tube,  in  the  form  of  a  sort  of  mirror. 
Apparently  the  vapor  of  some  metal  is  coming  off  with  the  oxygen 
and  condensing  on  the  cool  parts  of  the  tube.  As  this  shining  sub- 
stance accumulates  it  takes  the  form  of  globules,  which  may  be 
scraped  together.  It  is,  in  fact,  the  metal  mercury,  or  quicksilver. 
If  the  heating  continues  long  enough,  the  whole  of  the  red  powder 
eventually  disappears,  and  is  converted  into  these  two  products. 

Second  Variety  of  Chemical  Change:    Decomposition.  — 

Priestley's  experiment  introduces  to  us  a  second,  and  very  common 


CHEMICAL   CHANGE   AND   METHODS   OF   STUDYING   IT       15 

kind  of  chemical  action.  The  first  variety  was  combination  or 
union  (p.  7).  The  second  is  called  decomposition.  It  consists  in 
starting  with  a  single  substance  (here  mercuric  oxide)  and  splitting 
it  into  two  (or  more)  substances,  which  differ  in  properties  from  the 
substance  taken  and  from  one  another.  Here,  the  red  powder 
gave  mercury,  a  liquid  metal,  and  oxygen,  a  colorless  gas. 

Simple  and  Compound  Substances.  —  We  have  seen  that 
two  (or  more)  substances,  like  lead  and  oxygen,  can  combine  to 
form  a  compound  substance.  Are  all  substances,  then,  com- 
pounds? We  find  that  some  are  not.  We  have  never  succeeded 
in  obtaining  lead,  or  oxygen,  or  iron,  or  tin,  or  sulphur  by  com- 
bining any  two  substances.  We  can  decompose  mercuric  oxide 
by  heat,  and  we  have  other  ways  of  decomposing  compounds  like 
oxide  of  tin  and  ferrous  sulphide,  but  we  have  never  succeeded  in 
decomposing  the  mercury  or  the  oxygen,  the  iron  or  the  sulphur 
themselves.  Substances  which  we  are  not  able,  at  will,  to  decompose 
into,  or  to  make  by  chemical  union  from,  other  substances  are  called 
simple  or  elementary  substances.^  The  distinction  between  simple 
and  compound  substances  was  first  drawn  by  Lomonossov  in  1741. 
Later,  and  independently,  it  was  stated  very  clearly  by  Lavoisier 
(1789). 

Several  substances,  regarded  in  Lavoisier's  time  as  elementary, 
have  since  been  shown  to  be  compounds.  Thus,  quicklime  was  a 
simple  substance  until  Davy,  in  1808,  prepared  the  metal  calcium 
and  showed  that  quicklime  was  the  oxide  of  this  metal.  Hence, 
we  do  not  say  that  the  substances  regarded  as  simple  cannot  be 
decomposed,  but  only  that  they  are  substances  which  we  "are  not 
able"  (at  present)  to  decompose. 

The  phrase  "at  will"  is  also  important.  Radium  (q.v.*)  cannot 
be  decomposed  at  will,  but  it  undergoes  continuous  "disintegra- 
tion," producing  the  elements  helium  and  lead.  We  can  neither 
hasten,  retard,  nor  stop  this  spontaneous  decomposition. 

The  highly  interesting  experiments  of  Collie,  Paterson,  and 
Masson  (Chemical  Society,  London,  Annual  Report,  1914,  pp.  41- 
47)  seem  to  show  that  the  elements  helium  and  neon  can  be  pro- 

*  Contraction  for  quod  vide,  which  see.  This  abbreviation  is  used  when 
subjects  not  yet  discussed  are  mentioned.  For  such  subjects,  consult  the 
index. 


16  COLLEGE   CHEMISTRY 

duced  by  electrical  discharges  in  vacuum  tubes  and  even  in  a 
closed  tube  surrounding  the  vacuum  tube.  Before  long,  there- 
fore, the  decomposition  of  elementary  substances,  and  the  forma- 
tion of  some  elements  from  others,  at  mil,  may  be  a  recognized 
possibility,  and  the  foregoing  definition  may  have  to  be  radically 
revised. 

Elements.  —  The  word  element  is  used  in  two  senses.  It  is 
applied  to  the  simple  substance.  Thus  we  speak  of  "the  element 
iron,"  meaning  the  metal  iron.  It  is  applied  also  to  the  iron- 
matter  contained  in  ferrous  sulphide  or  in  ferric  oxide.  The 
reader  should  note  that  it  is  correct  usage  to  speak  of  the  element 
iron  and  the  element  sulphur  in  ferrous  sulphide,  but  a  chemist 
would  never  say  that  this  compound  contained  the  simple  sub- 
stances iron  and  sulphur.  If  he  did,  we  should  understand  him  to 
mean  that  it  was  a  mixture,  and  we  should  expect  parts  of  the 
material  to  be  magnetic  like  iron,  and  other  parts  to  be  yellow  and 
soluble  in  carbon  disulphide,  which  is  not  the  case.  In  the  same 
way  the  name  of  an  element  (such  as  iron)  is  applied  both  to 
the  material  in  combination  and  to  the  free  substance.  Thus 
"iron"  may  mean  free,  uncombined,  metallic  iron,  or  iron-matter 
in  some  compound.  The  sense  in  which  the  word  is  employed 
must  be  inferred  from  the  context  or  circumstances.  When  a 
chemist  speaks,  as  he  sometimes  does,  colloquially,  of  "iron"  in  a 
drinking  water,  for  example,  we  know  at  once  that  he  refers  to 
iron  in  the  form  of  some  compound,  for  metallic  iron  does  not 
dissolve  in  water  and,  if  it  did,  would  quickly  turn  into  rust  or 
some  other  form  of  combination. 

The  word  element,  then,  means  one  of  the  simple  forms  of 
matter,  either  free  or  in  combination. 

In  formally  describing  a  body  or  specimen,  the  chemist  always 
avoids  the  ambiguity  just  referred  to  by  naming  the  components, 
i.e.,  the  substance  or  substances  it  contains.  He  assumes  that  the 
nature  and  constituents  of  these  substances  will  be  known  to  any- 
one hearing  or  reading  the  description.  If  he  says  the  body  con- 
tains zinc  and  sulphur,  it  is  understood  that  the  body  is  a  mixture 
of  these  simple  substances.  If  it  contained  these  elements  in 
combination,  the  chemist  would  report  that  it  was  sulphide  of 
zinc. 


CHEMICAL   CHANGE    AND    METHODS    OF   STUDYING    IT      17 

The  Common  Elements.  —  Thousands  of  different  com- 
pound substances  are  known  but,  when  they  are  decomposed,  it  is 
found  that  the  number  of  different  elements  contained  in  them  is 
not  great.  Dozens  of  substances  contain  iron,  hundreds  contain 
sulphur,  thousands  contain  oxygen.  In  fact,  by  combining  a 
limited  number  of  simple  substances,  two,  three,  or  four,  together, 
in  varying  proportions  by  weight,  an  almost  unlimited  number  of 
different  compound  substances  could  be  produced. 

A  list  of  the  elements  appears  on  the  inside  of  the  cover,  at  the 
end  of  this  book,  and  contains  about  eighty  names.  Of  these,  a 
large  number  are  rare,  and  seldom  encountered.  More  than  99 
per  cent  of  terrestrial  material  is  made  up  of  eighteen  or  twenty 
elements  and  their  compounds.  Only  about  twenty  elements 
occur  in  nature  in  their  simple,  uncombined  condition.  Three- 
fourths  of  the  whole  number  are  found  in  combination  exclusively, 
and  must  be  liberated  by  some  chemical  action. 

Taking  the  atmosphere,  all  terrestrial  waters,  and  the  earth's 
crust,  so  far  as  it  has  been  examined,  F.  W.  Clarke  has  estimated 
the  plentifulness  of  the  various  elements.  The  first  twelve,  with 
the  quantity  of  each  contained  in  one  hundred  parts  of  terrestrial 
matter,  and  constituting  together  99  per  cent,  are  as  follows: 

Oxygen.    .    .  49.85  Calcium     ....  3.18  Hydrogen  .  .  0.97 

Silicon    .    .    .  26.03  Sodium 2.33  Titanium  .  .  0.41 

Aluminium    .     7.28  Potassium.    ...  2.33  Chlorine  .  .  .0.20 

Iron    ....     4.12  Magnesium  .    .    .  2.11  Carbon    .  .  .  0.19 

Thus  oxygen  accounts  for  nearly  one-half  of  the  whole  mass. 
Silicon,  the  oxide  of  which  when  pure  is  quartz  and  in  less  pure 
form  constitutes  ordinary  sand,  makes  up  half  of  the  remainder. 
Valuable  and  useful  elements,  like  gold,  silver,  sulphur,  and  mer- 
cury, are  among  the  less  plentiful  which,  all  taken  together,  furnish 
the  remaining  one  per  cent. 

Law  of  Definite  Proportions.  — In  the  decomposition  of 
mercuric  oxide  (p.  14)  we  find  that,  for  every  100  parts  of  mercury 
liberated,  almost  8  parts  of  oxygen  (more  exactly,  7.97  parts)  by 
weight  are  set  free.  Using  the  numbers  commonly  employed  in 
chemistry,  which  represent  the  same  proportion  by  weight: 
Mercuric  oxide  (216.6)  ->  Mercury  (200.6)  +  Oxygen  (16). 


18  COLLEGE    CHEMISTRY 

We  find  also  that  mercury  and  oxygen  can  be  made  to  combine  to 
form  mercuric  oxide,  and  the  proportions  by  weight  required  are 
the  same.  Moreover,  every  sample  of  mercuric  oxide,  whether 
made  by  combination,  or  in  any  of  the  other  possible  ways,  always 
contains  this  proportion  of  the  two  elements.  We  have  already 
seen  that  the  oxides  of  lead  and  tin  contain  fixed  proportions  (p.  9) 
of  the  metal  and  oxygen  and  that  ferrous  sulphide  has  a  constant 
composition  by  weight.  The  same  principle  is  found  to  apply  to 
all  chemical  compounds,  and  is  stated  in  the  law  of  definite  or 
constant  proportions:  In  every  sample  of  any  compound  substance, 
formed  or  decomposed,  the  proportion  by  weight  of  the  constituent 
elements  is  always  the  same.  (For  the  only  known  exception  to 
this  law,  see  radium.) 

Conservation  of  Mass.  —  The  most  painstaking  chemical 
work  seems  to  show  that,  if  all  the  substances  concerned  in  a 
chemical  change  are  weighed  before  and  after  the  change,  there  is 
no  evidence  of  any  alteration  in  the  quantity  of  matter.  The  two 
weights,  representing  the  sums  of  the  constituents  and  of  the  prod- 
ucts, respectively,  are,  indeed,  never  absolutely  identical,  but 
the  more  careful  the  work  and  the  more  delicate  the  instrument 
used  in  weighing,  the  more  nearly  do  the  values  approach  identity. 
We  are  able  to  state,  therefore,  that  the  mass  of  a  system  is  not 
affected  by  any  chemical  change  within  the  system. 

This  statement  simply  means  that  the  great  law  of  the  conserva- 
tion of  mass  holds  true  in  chemistry  as  it  does  in  physics.  Chemi- 
cal changes,  thoroughgoing  as  they  are  in  respect  to  all  other 
qualities,  do  not  affect  the  mass;  an  element  carries  with  it  its 
weight,  entirely  unchanged,  through  the  most  complicated  chemi- 
cal transformations. 

Superficial  observation,  as  of  a  growing  tree,  might  seem  to  give 
evidence  of  the  very  opposite  of  conservation  of  matter.  But 
here  the  carbon  dioxide  gas  in  the  air,  the  most  important  source 
of  nourishment  for  plants,  is  overlooked.  Similarly,  the  gradual 
disappearance  of  a  candle  by  combustion  seems  to  illustrate  the 
destruction  of  matter.  But  if  we  catch  the  gases  which  rise  through 
the  flame  (Fig.  12),  we  find  that  the  gases  weigh  even  more  than 
the  part  of  the  candle  which  has  been  sacrificed  in  making  them. 
When  we  take  account  of  the  weight  of  the  oxygen  obtained  from 


CHEMICAL   CHANGE    AND   METHODS   OF   STUDYING   IT      19 


the  air  which  sustains  the  combustion,  we  find  that  there  is  really 
neither  loss  nor  gain  in  weight.  If  we  carry  out  chemical  changes 
in  closed  vessels  (Fig.  13),  which  permit  neither  escape  nor  access 
of  material,  we  find  that  the  weight  does  not  alter. 


Fia.  12. 


Fia  13. 


Specific  Physical  Properties.  —  It  will  be  seen  that,  to  the 
chemist,  knowing  the  physical  properties  of  all  substances  is  very 
important.  By  means  of  the  properties,  he  recognizes  and  de- 
scribes all  the  bodies  he  studies.  It  may  be  well,  therefore,  here 
to  give  a  list  of  the  more  important  properties,  most  of  which  have 
been  mentioned  in  connection  with  the  illustrations  we  have  used. 

In  the  case  of  solids,  the  chief  physical  properties  the  chemist 
uses  are  color,  crystalline  form,  solubility  or  non-solubility  in  water 
and  occasionally  other  liquids,  the  temperature  at  which  the  sub- 
stance melts  (melting-point),  and  the  density. 

In  the  case  of  liquids,  he  notes  the  temperature  at  which  the 
liquid  boils  (boiling-point),  the  density,  the  mobility,  the  odor,  and 
the  color. 

Finally,  in  the  case  of  gases,  the  properties  commonly  mentioned 
are  the  color,  taste,  and  odor,  the  density,  solubility  in  water,  and 
the  ease  or  difficulty  with  which  the  gas  can  be  liquefied. 

Attributes  and  Conditions.  —  There  are  other  qualities 
which  a  body  may  possess  that  we  are  liable  to  confuse  with  the 
specific  properties.  Thus,  the  weight  of  a  piece  of  sulphur  is  not 


20  COLLEGE    CHEMISTRY 

a  property  of  sulphur.  A  hundred  pieces  of  as  many  different 
substances  might  all  have  the  same  weight,  so  that  a  particular 
weight  (say  2  grams)  is  not  a  property  of  any  one  species  of  matter. 
Weight,  dimensions,  and  volume  are  attributes  of  a  body.  They 
have  different  values  for  different  bodies,  even  when  those  bodies 
are  all  composed  of  the  same  substance.  The  attributes  are 
physical  in  nature.  They  are  of  great  importance  in  chemistry, 
however,  because  they  afford  the  only  means  we  have  of  measuring 
quantities  of  substances. 

There  are  still  other  qualities  which  a  body  (or  specimen  of 
matter)  may  possess.  It  has,  for  example,  a  certain  temperature, 
pressure  (state  of  compression),  motion,  or  electric  charge,  and  it 
may  be  in  solution  in  some  liquid.  A  body  may  change  in  tem- 
perature, pressure,  or  state  of  electrification,  or  it  may  be  dissolved 
in  water,  or  be  recovered  by  evaporation  of  the  liquid,  and  yet  be 
the  same  specimen.  A  hundred  specimens  of  as  many  different 
substances  may  all  have  the  same  temperature  —  this  is  not  a 
specific  property.  These  are  all  spoken  of  as  conditions.  They 
are  physical  conditions.  In  chemistry,  conditions  are  often  altered 
in  order  to  bring  about,  or  to  stop  chemical  change,  or  to  modify 
the  speed  with  which  it  takes  place.  Thus,  we  heated  the  lead 
(raised  its  temperature)  in  order  to  hasten  the  process  of  rusting. 
If  a  substance,  or  mixture,  is  capable  of  undergoing  ordinary 
chemical  change,  then  the  change  is  always  hastened  by  raising 
the  temperature,  and  is  always  delayed  or  prevented  by  lowering 
the  temperature.  Similarly,  changing  the  pressure  in  a  gas,  or 
dissolving  a  substance  in  some  liquid,  frequently  hastens  or  de- 
lays a  chemical  change  in  which  the  substance  takes  part.  The 
proper  physical  conditions  are,  therefore,  considered  in  connection 
with  every  chemical  operation.  Conditions  are  used  to  modify 
chemical  change. 

Physics  in  Chemistry.  —  It  will  be  seen  that  one  cannot 
accomplish  anything  in  chemistry  without  acquiring  and  using 
some  knowledge  of  physics.  We  measure  quantities  by  means  of 
the  physical  attributes,  weight  and  volume.  We  produce  chemi- 
cal change  by  arranging  the  physical  conditions,  for  example,  by 
mixing,  heating,  or  using  an  electric  current.  Physical  means  are 
the  only  means  we  possess  for  producing,  stopping,  or  modifying 


CHEMICAL   CHANGE   AND   METHODS   OF   STUDYING   IT      21 

chemical  changes.  Again,  we  ascertain  whether  a  chemical  change 
has  taken  place  or  not  by  observing  the  physical  properties  of 
the  materials  before  and  after  the  experiment.  Thus,  we  noted 
that  the  red,  powdery  oxide  of  mercury,  when  heated,  gave  a 
liquid  metal  and  a  gas.  All  the  phenomena  of  chemistry  are 
physical.  A  phenomenon  is  literally  something  that  is  seen  or, 
more  generally,  something  that  affects  any  of  the  senses.  Observ- 
ing physical  phenomena  is,  therefore,  our  sole  means  of  studying 
chemical  changes.  Chemical  work  is,  in  fact,  entirely  dependent 
upon  the  skilful  use  of  physical  agencies,  and  upon  the  close  obser' 
vation  of  physical  phenomena  for  its  success. 

It  is  only  the  inference,  following  the  experiment  and  the  obser- 
vation, that  is  strictly  chemical.  If  one  substance  gives  two 
different  substances,  or  if  two  substances  give  one  different  sub- 
stance, for  example,  we  infer  that  a  chemical  change  has  occurred. 
We  then  try  to  recognize  the  substances  by  their  properties  and 
name  them. 

Changes  like  that  of  ice  into  water,  or  of  water  into  steam,  and 
vice  versa,  are  not  regarded  as  chemical  changes.  These  are  called 
changes  of  state,  or  of  state  of  aggregation.  The  solid,  liquid,  and 
gaseous  forms  are  different  states  of  the  same  substance. 

Law:  Explanation:  Scientific  Method.  —  There  is  a 
widely  spread  impression  that  a  science,  like  chemistry,  is  a  part  of 
the  natural  order  of  the  universe.  It  is  thought  that  we  are  try- 
ing to  find  the  boundaries  of  chemistry,  as  they  have  been  pre- 
determined by  nature,  and  to  discover  the  facts,  relations  of 
facts,  and  laws  which  nature  has  provided  as  a  means  of  classifying 
the  content  of  the  science.  Now,  the  situation  is  precisely  the 
reverse  of  this.  Nature  provides  only  the  materials  and  the  phe- 
nomena, and  man  is  attempting  to  classify  them.  He  divides  the 
whole  into  groups,  such  as  physics,  chemistry,  botany,  etc.  Then 
he  classifies  the  facts  within  each  group,  in  order  that  he  may 
more  easily  remember  them  and  perceive  their  relations.  He 
often  finds  that,  when  new  facts  are  discovered,  parts  of  the 
classification  have  to  be  changed.  Thus,  as  we  have  just  seen, 
changes  of  state  are  usually  assigned  to  physics,  but  Ostwald  at 
one  time  suggested  that  they  should  be  considered  as  chemical 
phenomena. 


22  COLLEGE    CHEMISTRY 

In  the  preceding  pages,  we  have  discussed  some  of  the  ways  that 
have  been  invented  for  classifying  the  materials  and  facts  assigned 
to  chemistry.  Thus,  we  pick  out  a  number  of  facts  of  a  like 
nature  and  try  to  make  a  single  statement  which  will  cover  all 
these  facts.  For  example,  we  find  about  one  hundred  thousand 
different  substances  and,  in  the  case  of  each  substance,  every  speci- 
men that  we  have  examined  contains  the  same  proportions  of 
the  constituent  elements.  So  we  formulate  the  law  of  constant 
proportions.  A  law  or  generalization  in  chemistry  is  a  brief  state- 
ment describing  some  general  fact  or  constant  mode  of  behavior. 
We  must  remember,  however,  that  laws  are  only  true  so  long  as  no 
facts  in  conflict  with  them  are  known.  There  are  no  laws  in  nature. 
Nature  presents  materials  and  phenomena  as  she  pleases.  The 
laws  are  parts  of  science,  which  is  made  by  man,  and  is  a  description 
of  natural  facts  as  man  knows  them.  As  we  have  seen  (p.  18), 
at  least  one  undoubted  exception  to  the  law  of  constant  propor- 
tions has  recently  (1914)  been  discovered,  and  other  exceptions  to 
this  law  will  undoubtedly  present  themselves. 

One  section  (p.  5)  was  entitled:  "Explanation  of  rusting." 
If  that  paragraph  be  now  re-read,  it  will  be  found  that,  in  the 
ordinary  (as  distinct  from  the  scientific)  sense  of  the  word,  no 
explanation  was  given!  When  we  ask  a  man  to  "explain"  some 
feature  in  his  conduct,  we  recognize  that  he  might  have  chosen  to 
act  otherwise,  and  we  wish  to  know  why  he  acted  precisely  as  he 
did.  Nature,  however,  has  no  free  will,  and  cannot  tell  why  she 
presents  certain  phenomena,  and  not  others. 

On  examining  the  explanation,  we  find  that  it  simply  shows 
that  when  iron  rusts  it  combines  with  oxygen  from  the  air.  This 
is  an  additional  fact.  It  shows  how  iron  rusts,  namely  by  taking 
up  oxygen,  but  not  why  it  is  able  to  unite  with  oxygen.  We 
simply  do  not  know  why  iron  can  combine  with  oxygen  gas  and 
platinum  cannot. 

Explanations  in  chemistry  are  of  three  kinds.  (1)  We  usually 
try  to  show  that  the  phenomenon  is  not  an  isolated  one.  Thus, 
we  show  that  other  metals  rust.  This  reconciles  us  to  some  ex- 
tent .to  the  fact  that  iron  rusts,  and  we  feel  some  mental  satisfac- 
tion. This  is  the  method  of  showing  that  the  fact  to  be  explained 
is  a  member  of  a  large  class  of  similar  facts.  (2)  Next,  we  try  to  get 
more  information  about  the  fact  to  be  explained.  Thus,  when,  to  the 


CHEMICAL   CHANGE   AND   METHODS   OF   STUDYING   IT      23 

acquaintance  with  the  outward  manifestations  of  rusting,  we  add 
the  further  information  that  there  is  an  increase  in  weight,  and 
that  this  is  due  to  union  of  oxygen  from  the  air  with  the  iron,  we 
feel  increased  satisfaction,  and  say  that  the  fact  has  been  "ex- 
plained." (3)  If  we  are  still  dissatisfied,  and  can  discover  no 
further  useful  facts,  we  imagine  a  state  of  affairs  which,  if  true, 
would  classify  the  fact  or  add  to  what  we  know  about  it.  This  step 
we  call  explaining  by  means  of  an  hypothesis.  We  then  devote 
our  attention  to  trying  to  verify  the  hypothesis. 

The  formulation  of  laws  and  the  making  of  attempts  to  explain 
facts  are  part  of  what  is  called  the  scientific  method.  The  purpose 
of  this  method  is  to  convert  the  subject  matter  into  a  science,  that 
is,  into  an  organized  body  of  knowledge. 

Summary.  —  In  this  chapter  we  have  learned:  (1)  that, 
while  there  are  many  substances,  there  is  a  limited  number  of 
entirely  different  kinds  of  matter  (elements) ;  (2)  that,  in  addition 
to  constant  physical  properties,  each  substance  has  a  constant 
composition  by  weight.  We  have  also  learned  that  physical 
properties  are  utilized  in  manipulations,  like  filtration  and  evap- 
oration, as  well  as  for  identifying  substances,  and  that  physical 
attributes  are  used  for  measuring  quantities  in  chemistry  and 
physical  conditions  for  guiding  chemical  change.  Finally,  we 
have  seen  that  a  science  is  not  a  natural,  but  a  manufactured 
product,  and  that  the  science  of  chemistry  is  still  in  the 
making. 

Exercises.* — 1.  What  physical  properties  are  used  (a)  in 
filtration,  (6)  in  evaporation,  (c)  in  the  separation  and  identifica- 
tion of  the  products  from  heating  mercuric  oxide  (p.  14)? 

2.  Describe:  (a)  a  red-hot  rod  of  iron,  10  cm.  long  by  1  cm. 
diameter,  weighing  58.5  g.;  (6)  a  solution  of  5  g.  of  sulphur  in 
20  c.c.  (59  g.)  of  carbon  disulphide  at  18°  C.  In  doing  so,  divide 
the  description  into  attributes,  conditions,  and  properties. 

*  The  exercises  should  in  all  cases  be  studied  with  minute  care.  They  not 
only  serve  as  tests  to  show  that  the  chapter  has  been  understood,  but  very 
frequently  (as  in  No.  4)  also  call  attention  to  ideas  which  might  not  be  ac- 
quired from  the  text  alone,  or  (as  in  Nos.  1,  2,  5)  assist  in  elucidating  ideas 
given  in  the  text  which,  without  the  exercises,  might  not  be  fully  grasped. 


24  COLLEGE    CHEMISTRY 

3.  Consider  the  following  materials  and  state  whether,  so  far 
as  you  can  now  judge,  each  is  a  single  substance  or  a  mixture:  (a) 
a  candle,  (b)  a  cake  of  soap,  (c)  an  egg. 

4.  What  are  the  two  most  direct  ways  of  showing  a  substance 
to  be  a  compound?    Illustrate  each. 

5.  If  we  say  that  quicklime  contains  calcium  (p.  15),  do  we 
mean  the  element  or  the  simple  substance  calcium? 

6.  What  explanation  was  given,  (a)  of  the  disappearance  of 
mercuric  oxide  when  heated,  (6)  of  the  absence  of  iron  and  sulphur, 
as  substances,  from  ferrous  sulphide?    Which  of  the  three  kinds 
of  explanation  was  used  in  each  case? 


CHAPTER  III 
OXYGEN 

WE  cannot  do  better  than  begin  the  more  systematic  study  of 
chemistry  with  oxygen,  for  it  is  a  most  interesting  as  well  as  useful 
substance.  It  is  the  active  component  of  the  air.  We  depend 
upon  it  for  life,  since  in  its  absence  we  suffocate,  for  heat,  since 
wood,  coal,  and  gas  will  not  burn  without  it,  and  even  for  light 
where  oil,  gas,  or  a  candle  is  used. 

We  wish  to  know  with  which  substances  we  use  in  the  laboratory 
it  can  combine,  as  well  as  the  substances  on  which  it  has  no 
action.  This  information  will  show  us  how  to  work,  in  future, 
without  interference  from  the  oxygen  in  the  air  and  whether  oxygen 
has  probably  played  a  part  in  some  experiment  or  not. 

Let  us  take  up,  then,  (1)  the  history  of  the  element,  (2)  what 
materials  contain  oxygen  (occurrence),  (3)  how  we  can  obtain  it  in 
a  pure  state  (preparation),  (4)  what  its  specific  physical  properties 
as  a  substance  are,  and  (5)  what  it  does,  and  what  it  cannot  do  in 
nature  and  in  the  laboratory  (chemical  properties).  The  classifi- 
cation of  the  facts  about  this,  and  other  substances,  under  five 
heads  is  somewhat  mechanical,  but  has  the  advantage  of  enabling 
the  reader  quickly  to  find  any  required  information. 

History  of  Oxygen.  —  The  Chinese,  in  or  before  the  eighth 
century,  knew  that  there  were  two  components  in  the  air,  and 
that  the  active  one,  yin,  combined  with  some  metals,  and  with 
burning  sulphur,  and  charcoal.  They  even  knew  that  it  could  be 
obtained  in  pure  form  by  heating  certain  minerals,  of  which  one 
was  saltpeter.  Leonardo  da  Vinci  (1452-1519)  seems  to  be  the 
first  European  to  mention  the  former  fact.  Mayow  (1669) 
measured  the  proportion  of  oxygen  in  the  air  and  discussed  fully 
its  uses  in  combustion,  rusting,  vinegar-making,  and  respiration, 
but  did  not  make  a  pure  sample.  Hales  (1731)  made  it  from  salt- 
peter, and  measured  the  amount  obtainable,  but  did  not  see  any 

25 


26  COLLEGE    CHEMISTRY 

connection  between  it  and  the  air!  Bayen  (Apr.,  1774)  was  the 
first  to  make  it  by  heating  mercuric  oxide.  Priestley  (Aug.  1, 
1774)  made  it  by  heating  the  same  substance  and  quite  purpose- 
lessly, as  he  admits,  thrust  a  lighted  candle  into  it  and  was  de- 
lighted with  the  extreme  brilliance  of  the  flame.  He  had,  however, 
entirely  incorrect  ideas  about  its  nature,  and  no  notion  until  a 
year  later  that  it  was  a  component  of  the  air.  Scheele,  a  Swedish 
apothecary,  had  made  it  in  1771-2  from  no  less  than  seven  different 
substances  and  understood  clearly  that  atmospheric  oxygen  com- 
bined with  metals,  phosphorus,  hydrogen,  linseed  oil  and  many 
other  substances.  But  the  publisher  did  not  get  his  book  out 
until  1777,  and  Priestley  is  usually  credited  with  the  "discovery" 
of  the  element.  Finally,  Lavoisier  (1777)  heated  mercury  in  a 
retort  (Fig.  14),  the  neck  of  which  projected  into  a  jar  standing 
in  a  larger  dish  of  mercury.  The  air,  thus  enclosed  within  the  jar 
and  the  retort,  during  twelve  days  lost 
one-fifth  of  its  volume.  Simultaneously, 
red  particles  of  mercuric  oxide  accumu- 
lated on  the  surface  of  the  mercury  in  the 
retort.  The  residual  gas,  nitrogen,  no 
longer  supported  life  or  combustion.  The 
oxide,  on  being  heated  more  strongly,  by 
itself,  gave  off  a  gas  whose  volume  exactly 

corresponded  with  the  shrinkage  undergone  by  the  enclosed  air, 
and  ;this  gas  possessed  in  an  exaggerated  degree  the  properties 
which  the  air  had  lost.  The  proof  that  oxygen  was  a  component 
of  the  atmosphere  was  therefore  complete.  Later,  Lavoisier,  in 
the  mistaken  belief  that  the  new  element  was  an  essential  con- 
stituent of  all  sour  substances,  named  it  oxygen  (Gk., acid-producer). 

Occurrence.  —  As  we  have  seen,  nearly  50  per  cent  of  terres- 
trial matter  is  oxygen.  Water  contains  about  89  per  cent,  the 
human  body  over  60  per  cent,  and  common  materials  like  sand- 
stone, limestone,  brick,  and  mortar  more  than  50  per  cent  of  this 
element.  One-fifth  by  volume  (nearly  one-fourth  by  weight)  of 
the  air  is  free  oxygen. 

Preparation  of  Oxygen.  —  1.   The  oxygen  of  commerce  is 
now  made  chiefly  from  liquefied  air  (q.v.*).     The  liquid  oxygen 
*  See  p.  15,  footnote. 


OXYGEN  27 

boils  at  — 182.4°,  but  the  nitrogen  boils  at  an  even  lower  tempera- 
ture (  —  194°).  Since  the  liquid  air  has  a  temperature  of  about 
—  190°,  somewhat  above  that  of  boiling  nitrogen,  the  latter 
evaporates  much  more  freely  than  does  the  oxygen.  After  a 
time,  when  the  remaining  liquid  is  almost  pure  oxygen  (96  per 
cent),  the  gas  coming  off  is  compressed  by  pumps  into  the  steel 
cylinders  (Fig.  15)  in  which  it  is  sold.  In  medicine,  patients 
suffering  from  pneumonia  or  suffocation  obtain  some  relief  by 
inhaling  it  in  this  form.  It  is  also  used  in  feeding 
flames,  instead  of  air,  when  intense  heat  is  required 
(see  acetylene  torch  and  calcium  light) . 

2.  Unfortunately,  it  is  difficult  to  liberate  oxygen 
from  natural  substances.  Saltpeter  (potassium  nitrate), 
for  example,  which  is  found  in  many  soils  and  can 
be  dissolved  out  with  water,  gives  off  oxygen  (p.  25) 
only  when  raised  to  a  bright  red  heat  by  the  Bunsen 
flame  or  blast  lamp.     But,  even  at  this  temperature, 
it  gives  up  only  one-third  of  the  oxygen  it  contains. 

3.  In  practice,  we  are  compelled  to  use  manufac- 
tured substances.     Amongst  the  artificial  substances 
are  mercuric  oxide,  expensive  but  historically  inter- 
esting (p.  14),  potassium  chlorate,  perhaps  the  most 
convenient  for  laboratory  use,  and  sodium  peroxide. 
Potassium  Chlorate   (q.v.)  is  a  white  crystalline  substance  used, 
on  account   of  the   oxygen  it   contains,   in  large  quantities   in 
the  manufacture  of  matches  and  fireworks.     When  heated  in  a 
tube  similar  to  that  in  Fig.  11,  it  first  melts  (351°)  and  then,  on 
being  more  strongly  heated,  it  effervesces  and  gives  off  a  very  large 
volume  of  oxygen.     Examination  shows  that  the  whole  of  the 
oxygen  it  contains  (39  per  cent)  can  be  driven  out.     The  white 
material  which  remains  after  the  heating  is  identical  with  the 
mineral  sylvite.     To  the  chemist  it  is  known  as  potassium  chloride. 
The  change,  together  with  the  weights  of  the  materials,  is  as  follows: 

Potassium  chlorate  (122.56)  -» Potassium  chloride  (74.56)  +  Oxygen  (48) 

Potassium  (39.1)  Potassium  (39.1) 

Chlorine  (35.46)  Chlorine  (35.46) 

Oxygen  (48) 

A  peculiarity  of  this  action  is  that  admixture  of  manganese 
dioxide    (the   mineral   pyrolusite)    increases   very  markedly  the 


FIG.  15. 


28 


COLLEGE   CHEMISTRY 


speed  with  which  the  decomposition  of  the  potassium  chlorate 
takes  place.  Hence,  in  its  presence,  and  it  is  generally  mixed 
with  the  chlorate  in  laboratory  experiments  (Fig.  16),  a  sufficient 
stream  of  the  gas  is  obtained  at  a  relatively  low  temperature 


Fio.  16. 


(below  200°,  see  p.  29).     Hales  (p.  25)  was  the  first  to  collect  a 

gas  over  water  (Fig.  16),  in  order  that  it  might  be  kept  unmixed 
with  air. 

4.  Oxygen  can  be  obtained  conveniently 
from  sodium  peroxide  and  water  by  means  of 
generators  (Fig.  17)  similar  to  the  acetylene 
generators  used  on  automobiles.  When 
the  metal  sodium  is  burned  in  air,  sodium 
peroxide  is  obtained  as  a  powder.  This 
powder,  after  being  melted,  solidifies  in  com- 
pact, solid  form,  and  is  sold  as  oxone.  The 
oxone  is  bought  in  a  small,  sealed  tin  can, 
the  ends  of  which  are  perforated  in  several 
places  just  before  use.  When  the  valve  (B) 
is  opened,  so  that  the  oxygen  escapes,  the 
water,  which  fills  the  generator  almost  to 
the  top,  enters  the  can  (C)  by  the  holes  in 
the  bottom  and  interacts  with  the  oxone. 
When  the  valve  is  shut,  the  gas  continues 
to  be  generated  until  it  has  driven  the 

water  down  again  below  the  level  of  the  bottom  of  the  can. 

Sodium  peroxide  (78)  +  Water  (18)  ->  Sodium  hydroxide  (80)  +  Oxygen  (16) 

Sodium  (46)  Hydrogen  (2.016)      Sodium  (46) 

Oxygen  (32)  Oxygen  (16)  Oxygen  (32) 

Hydrogen  (2.016) 


FIG.  17. 


OXYGEN 


29 


This  method  is  convenient  because  it  works  at  room  temperature 
and  can  be  started  and  stopped  at  will.  The  sodium  hydroxide 
produced  is  very  soluble  in  water  and  remains  dissolved.  Note 
that  the  name  of  this  substance  indicates  the  elements  which 
compose  it. 

Catalytic  or  Contact  Action.  —  The  influence  of  manganese 
dioxide  in  causing  the  potassium  chlorate  to  decompose  more  easily 
(p.  27)  well  deserves  notice.  The  effect  is  very  striking  if  some 
pure  potassium  chlorate  is  melted  carefully,  to  avoid  superheat- 
ing, in  a  wide-mouth  flask  (Fig.  18).  The  flask  is  provided  with 
a  wide  exit  tube,  from  which  a 
rubber  tube  may  lead  to  a  bottle 
inverted  in  a  trough  filled  with 
water  as  in  Fig.  16.  A  little 
manganese  dioxide  is  contained 
in  the  upper,  closed  tube.  No 
effervescence  of  the  chlorate 
can  be  seen  at  its  melting-point 
(334°)  — only  a  little  air,  ex- 
panded by  the  heating,  issues 
from  the  tube.  When,  however, 
the  closed  tube  containing  the 
manganese  dioxide  is  rotated 
into  a  vertical  position  (see 
dotted  lines) ,  and  the  black  powder  falls  into  the  chlorate,  the  oxygen 
comes  off  in  torrents,  in  consequence  of  the  enormous  acceleration  of 
the  decomposition.  As  a  precaution  against  injury  from  an  explosion, 
it  is  advisable  to  wrap  the  flask  in  a  towel  before  turning  the  tube. 

It  must  also  be  noted  that  the  manganese  dioxide  is  not  itself 
permanently  altered.  If  the  material  left  after  the  action  is 
shaken  with  water,  the  potassium  chloride  dissolves,  while  the 
dioxide  does  not.  Filtration  (p.  12)  then  enables  us  to  recover 
the  latter,  and  to  ascertain  that  it  has  been  changed  neither  in 
quantity  nor  in  properties. 

The  only  effect  of  the  dioxide  is  to  hasten  the  decomposition  of 
the  chlorate,  which  would  otherwise  be  too  slow  at  200°  (p.  28), 
or  even  at  334°  (its  m.-p.)  to  be  of  any  practical  value.  Sub- 
stances which  hasten  a  chemical  action  without  themselves  under- 


FIG.  18. 


30  COLLEGE    CHEMISTRY 

going  any  permanent  change  are  called  contact  agents,  catalytic 
agents,  or  catalysts.  The  process  is  called  contact  action  or  catal- 
ysis (Gk.,  decomposition,  not  a  very  fortunate  choice  of  words). 
Such  substances  are  frequently  used  in  chemistry.  The  addition 
of  a  suitable  catalyst  is  one  of  the  conditions  (p.  20)  for  carrying 
out  actions  in  which  a  contact  agent  is  necessary.  Many  sub- 
stances of  this  class  are  secreted  by  animals  and  plants  and  play 
an  important  part  in  digestion,  fermentation,  and  other  physiologi- 
cal changes.  Their  presence  often  enables  very  complex  chemi- 
cal actions  to  proceed  rapidly  at  rather  low  temperatures. 

The  oxone,  mentioned  above,  always  contains  a  trace  of  cuprous 
oxide  which  hastens  the  action  on  water. 

Specific  Properties  of  Two  Kinds,  Physical  and  Chemical. 

—  We  have  learned  that'every  substance  has  its  own  set  of  specific 
properties.  In  describing  a  substance,  it  is  convenient  to  divide  the 
properties  into  two  classes.  The  list  of  substances  with  which  the 
given  substance  can  enter  into  chemical  combination,  for  example, 
we  place  under  specific  chemical  properties.  Relations  of  the  sub- 
stance to  any  of  the  varieties  of  chemical  change  belong  to  this  class. 

On  the  other  hand,  we  do  not  consider  melting  or  boiling  to  be 
chemical  changes,  so  we  place  the  temperatures  at  which  the  sub- 
stance melts  (m.-p.)  and  boils  (b.-p.),  its  color,  etc.  (for  list,  see 
p.  19),  under  specific  physical  properties. 

Properties  of  either  class  may  be  used  for  recognizing  a  substance. 

Specific  Physical  Properties  of  Oxygen.  —  Oxygen  resembles 
air  in  having  neither  color,  taste,  nor  odor.  The  density  of  a  sub- 
stance is,  strictly  speaking,  the  weight  of  1  cubic  centimeter  (1  c.c.). 
In  the  case  of  a  gas,  we  frequently  prefer  to  give  the  weight  of  1000 
c.c.  (1  liter),  at  0°  and  760  mm.  (1  atmosphere)  barometric  pressure. 
For  oxygen  this  weight  is  1.42900  grams  (Morley).  The  corre- 
sponding weight  for  air  is  1.293,  so  that  oxygen  is  slightly  heavier, 
bulk  for  bulk,  than  air  (in  the  ratio  1.105  :  1).  Oxygen  can  be 
liquefied  by  compression,  provided  its  temperature  is  first  reduced 
below  —118°,  which  is  its  critical  temperature.*  The  gas  is 

*  Each  gas  has  an  individual  critical  temperature  (q.v.)  above  which  no 
pressure,  however  great,  will  produce  liquefaction.  The  farther  the  tempera- 
ture of  a  specimen  of  the  gas  is  below  the  critical  point,  the  less  will  be  the 
pressure  required  to  liquefy  it. 


OXYGEN 


31 


slightly  soluble  in  water,  the  solubility  at  0°  being  4  volumes  of  gas 
in  100  volumes  of  water  (at  20°,  3  :  100). 

The  solubility  of  oxygen  in  water,  although  slight,  is  in  some 
respects  its  most  important  physical  property.  Fish  obtain  oxy- 
gen for  their  blood  from  that  dissolved  in  the  water.  With  air- 
breathing  animals  (like  man),  the  oxygen  could  not  be  taken  into 
the  system,  if  it  did  not  first  dissolve  in  the  moisture  contained  in 
the  walls  of  the  air  sacs  of  the  lungs,  and  then  pass  inwards  in  a 
dissolved  state  to  the  blood. 

Liquid  oxygen,  first  prepared  by  Wroblevski,  has  a  pale-blue 
color.  At  one  atmosphere  pressure,  that  is,  in  an  open  vessel,  it 
boils  at  - 182.5°.  Its  density  (weight  of  1  c.c.)  is  1.13,  so  that  it 
is  slightly  denser  than  water.  By  cooling  with  a  jet  of  liquid 
hydrogen,  Dewar  froze  the  liquid  to  a  snow- 
like,  pale-blue  solid.  A  tube  of  liquid  oxygen 
is  noticeably  attracted  by  a  magnet. 

Six  Specific  Physical  Properties  of  Each 
Gas.  —  Although  every  substance  has  many 
physical  properties,  we  shall  mention  only 
those  which  are  used  in  chemical  work,  with 
occasionally  the  addition  of  any  peculiar  or 
unexpected  quality.  It  will  aid  the  memory 
to  recall  the  physical  properties  of  a  gas,  if 
we  note  that,  as  a  rule,  only  six  such  proper- 
ties are  mentioned:  (1)  color,  (2)  taste,  (3) 
odor,  (4)  density,  (5)  liquefiability,  defined  by 
the  critical  temperature,  (6)  solubility,  usually 
in  water  only. 

Specific  Chemical  Properties  of  Oxy 

gen.  —  The  chemical  properties  of  pure 
oxygen  are  like  those  of  atmospheric  air, 
only  more  pronounced. 

Non-metallic  Elements.  Sulphur,  when  raised  in  advance  to  the 
temperature  necessary  to  start  the  action,  unites  vigorously  with 
oxygen  (Fig.  19),  giving  out  much  heat  and  producing  a  familiar 
gas  having  a  pungent  odor  (sulphur  dioxide).  This  odor  is  fre- 
quently spoken  of  as  the  "smell  of  sulphur,"  but  in  reality  sulphur 


FIG.  19. 


32  COLLEGE    CHEMISTRY 

itself  has  no  odor,  and  neither  has  oxygen.  The  odor  is  a  property 
of  the  compound  of  the  two.  The  mode  of  experimentation  can 
be  changed  and  the  oxygen  led  into  sulphur  vapor  through  a  tube. 
The  oxygen  then  appears  to  burn  with  a  bright  flame,  giving  the 
same  product  as  before. 

Phosphorus,  when  set  on  fire,  blazes  in  oxygen  very  vigorously, 
forming  a  white,  powdery,  solid  oxide  —  phosphorus  pent  oxide. 
Burning  carbon,  in  the  form  of  charcoal  or  hard  coal,  glows  bril- 
liantly and  is  soon  burnt  up.  It  leaves  an  invisible,  odorless  gas  - 
carbon  dioxide.  At  high  temperatures,  oxygen  combines  readily 
with  one  or  two  other  non-metals  (e.g.,  silicon,  boron,  and  arsenic), 
and  to  a  small  extent  (1  per  cent  at  1900°)  with  nitrogen.  It  will 
not  combine  directly  with  chlorine,  bromine,  or ,  iodine,  although 
oxides  of  the  first  and  last  can  be  prepared  by  using  other  varieties 
of  chemical  change.  With  the  six  members  of  the  helium  family 
(q.v.),  of  which  no  compounds  are  known,  and  with  fluorine,  oxygen 
forms  no  compounds. 

Sulphur  (32.06) + Oxygen  (32) ^Sulphur  dioxide  (64.06). 
Phosphorus  (62. 08)+ Oxygen  (80)  — ^Phosphorus pentoxide  (142.08). 
Carbon  (12)+ Oxygen  (32) -»Carbon  dioxide  (44). 

Metallic  Elements.  Iron,  as  we  have  seen,  rusts  exceedingly 
slowly  in  air  and,  even  when  red-hot,  gives  hammer-scale,  the  black 
solid  which  is  broken  off  on  the  anvil,  rather  deliberately.  In  pure 
oxygen,  a  bundle  of  picture-wire,  if  once  ignited,  will  burn  with 
surprising  brilliancy,  throwing  off  sparkling  globules  of  the  oxide, 
melted  by  the  heat.  This  oxide  is  a  black,  brittle  substance, 
identical  with  hammer-scale,  and  different  from  rust  (ferric  oxide) . 
It  contains,  in  fact,  a  smaller  proportion  of  oxygen  than  does  the 
latter,  and  is  called  magnetic  oxide  of  iron. 

Iron  (167.52)  +  Oxygen  (64)  ->  Magnetic  oxide  of  iron  (231.52). 

All  the  familiar  metals,  excepting  gold,  silver,  and  platinum, 
when  heated,  combine  with  oxygen,  some  more  vigorously,  others 
less  vigorously  than  does  iron.  Oxides  of  the  three  metals  just 
named  can  also  be  made,  but  only  by  varieties  of  chemical  change 
other  than  direct  combination. 

Compound  substances,  if  they  are  composed  largely  or  entirely 
of  elements  which  combine  with  oxygen,  are  able  themselves  to 


OXYGEN  33 

interact  with  oxygen.  Usually,  they  produce  a  mixture  of  the 
same  oxides  which  each  element,  separately,  would  give.  Hence, 
wood,  which  is  composed  of  carbon  and  hydrogen  with  some 
oxygen,  when  burnt  in  oxygen,  produces  carbon  dioxide  and  water 
(oxide  of  hydrogen)  in  the  form  of  vapor.  Again,  carbon  disul- 
phide  burns  readily,  giving  carbon  dioxide  and  sulphur  dioxide, 
just  as  do  carbon  and  sulphur,  separately.  'Ferrous  sulphide  gives, 
similarly,  sulphur  dioxide  and  magnetic  oxide  of  iron. 

Tests.     A  Test  for   Oxygen.  —  A  test  is  a  property  which, 

because  it  is  easily  recognized  (a  strong  color,  for  example),  or  for 
some  other  sufficient  reason,  is  commonly  employed  in  recognizing 
a  substance. 

Oxygen,  as  we  have  seen  (p.  14),  when  pure,  is  recognized  by 
the  fact  that  a  splinter  of  wood,  glowing  at  one  end,  bursts  into 
flame  when  introduced  into  the  gas.  Only  one  other  gas  (see 
nitrous  oxide)  behaves  similarly. 

The    Measurement    of    Combining    Proportions.  —  In    a 

number  of  condensed  statements  we  have  given  the  proportions 
by  weight  of  the  materials  combining.  It  is  now  desirable  that 
we  should  know  how  the  necessary  measurements  are  made.  The 
most  exact  measurement  of  the  proportions  in  which  the  elements 
combine  to  form  compounds  involves  manipulations  too  elaborate 
to  be  gone  into  here.  One  or  two  brief  statements,  diagrammatic 
rather  than  accurate,  will  show  the  principles,  however. 

If  we  take  a  weighed 
quantity  of  iron  in  a  test- 
tube  and  heat  it  with  more 
than  enough  sulphur  (an 
excess  of  sulphur),  we  get 
free  sulphur  along  with  the 
ferrous  sulphide  (pp.  13-14), 
and  no  free  iron  survives. 
We  may  remove  the  free 
sulphur  by  washing  the  solid  FlG-  2a 

with  carbon  disulphide.  The  difference  between  the  weights  of  the 
ferrous  sulphide  and  the  iron  gives  the  amount  of  sulphur  combined 
with  the  known  quantity  of  the  latter. 


34  COLLEGE   CHEMISTRY 

As  an  example  of  the  study  of  the  combination  of  a  metal  with 
oxygen,  we  may  weigh  a  small  amount  of  copper  in  the  form  of 
powder  in  a  porcelain  boat  and  pass  oxygen  over  the  heated  metal 
(Fig.  20).  If  we  limit  the  oxygen,  part  of  the  copper  may  remain 
unaltered;  if  we  use  it  freely,  the  excess  will  pass  on  unchanged. 
The  original  weight  of  the  copper,  and  the  increase  in  weight, 
representing  oxygen,  give  us  the  data  for  determining  the  compo- 
sition of  cupric  oxide.  The  data  furnished  by  one  rough  lecture- 
experiment,  for  example,  were  as  follows: 

Weight  of  boat  +  copper 4.278  g. 

Weight  of  boat  empty    ....„*.   t 3.428  g. 

Difference  =  weight  of  copper    .    .   . 0.860g. 

Weight  after  addition  of  oxygen 4 . 488  g. 

Weight  without  oxygen 4.278  g. 

Difference  =  weight  of  oxygen 0.210g. 

The  proportion  of  copper  to  oxygen,  so  far  as  this  one  measure- 
ment goes,  is  therefore  85  :  21. 

The  results  of  quantitative  experiments  are  often  recorded  in  the 
form  of  parts  in  one  hundred.  To  find  the  percentage  of  each  con- 
stituent, we  observe  that  the  proportion  of  copper  is  85  :  85  -f  21, 
or  TW  of  the  whole.  That  of  the  oxygen  is  -ft*j  of  the  whole.  Thus 
the  percentages  are: 

Copper,  106  :  85  :  :  100  :  x.        x   =  80.2. 
Oxygen,  106  :  21  :  :  100  :  x1 '.      x'  =  19.8. 

Naturally,  the  mean  of  the  results  of  a  number  of  more  carefully 
managed^experiments  will  be  nearer  the  true  proportion.  The  per- 
centages at  present  accepted  as  most  accurate  are  79.9  and  20.1. 

In  the  case  of  mercuric  oxide,  we  may  decompose  a  known  weight 
of  the  oxide  (p.  14),  collect  the  mercury  and  weigh  it,  and  ascer- 
tain the  oxygen  by  difference. 

The  names  of  the  constituent  elements  in  a  compound,  together 
with  the  proportion  by  weight  in  which  they  are  present,  are  called 
the  composition  of  the  substance.  Thus,  the  composition  of  cupric 
oxide  is  copper  :  oxygen  :  :  79.9  :  20.1.  This  is  the  percentage  com- 
position, but  other  numbers  expressing  the  same  proportion  (such 
as  63.57  :  16)  will  serve  the  purpose. 

All  experiments  involving  measurement,  such  as  those  used  in 
determining  composition,  are  called  quantitative  experiments. 


OXYGEN  35 

Another  Quantitative  Experiment.  —  The  following  will 
show  how  the  combining  proportions  may  be  measured  when  the 
product  is  a  gas,  the  weight  of  which  must  be  ascertained.  Sul- 
phur burns  in  oxygen  to  form  sulphur  dioxide.  A  known  weight 
of  sulphur  is  placed  in  a  porcelain  boat  (Fig.  21),  which  has  already 
been  weighed.  The  U-shaped  tube  to  the  right  contains  a  solu- 
tion of  potassium  hydroxide,  which 
is  capable  of  absorbing  the  resulting 
gas.  The  oxygen  enters  from  the  left. 
When  the  sulphur  is  heated,  it  burns 
in  the  oxygen,  and  the  loss  in  weight 
which  the  boat  undergoes  shows  the  FlG<  2L 

amount  of  sulphur  consumed.  The  gain  in  weight  of  the  U-tube 
shows  the  weight  of  the  compound  produced.  By  subtracting, 
we  get  the  quantity  of  oxygen. 

In  one  experiment,  the  loss  in  weight  of  the  boat  and  its  con- 
tents (=  sulphur)  was  1.21  g.  The  weight  gained  by  the  U-tube 
was  2.42  g.  The  difference  (  =  oxygen)  is  1.21.  The  proportion 
of  sulphur  to  oxygen  in  sulphur  dioxide  is  therefore  1.21  :  1.21  or 
1  :  1  or,  in  percentages,  50  :  50.  This  proportion  is  very  close  to 
the  accepted  value  (p.  32),  32.06  :  32. 

The  same  method  could  be  used  for  carbon,  for  the  carbon 
dioxide  produced  would  be  absorbed  in  the  solution  of  potassium 
hydroxide. 

Combustion.  —  Violent  union  with  oxygen  is  called,  in 
popular  language,  combustion  or  burning.  Yet,  since  oxygen  is 
only  one  of  many  gaseous  substances  known  to  the  chemist,  and 
similar  vigorous  interactions  with  these  gases  are  common,  the 
term  has  no  scientific  significance.  Even  the  union  of  iron  and  sul- 
phur gives  out  light  and  heat,  and  is  quite  similar  in  the  chemical 
point  of  view  to  combustion. 

A  misleading  term  often  used  in  this  connection  is  kindling 
temperature.  It  gives  the  impression  that  there  is  a  definite  tem- 
perature at  which  combustion  will  start.  But  the  temperature 
is  only  one  of  the  conditions  which  produce  combustion.  Finely 
powdered  iron  will  start  burning  at  a  lower  temperature  than  will 
an  iron  wire,  because  it  presents  relatively  more  surface  to  the  gas. 
Again,  if  the  oxygen  is  at  less  than  one  atmosphere  pressure,  the 


36  COLLEGE    CHEMISTRY 

wire  will  require  to  reach  a  higher  temperature  before  combustion 
will  begin.  Finally,  the  vapor  of  methyl  alcohol  and  air  requires 
to  be  raised  above  a  red  heat  before  combustion  starts,  but  a  pocket 
cigar-lighter  sets  fire  to  this  very  mixture  by  means  of  a  contact 
agent  (a  thin  platinum  wire)  without  any  other  means  of  heating 
being  required.  So  that,  the  conditions  under  which  combustion 
begins  involve  the  physical  condition  of  the  solid,  the  pressure  of 
the  gas  or  vapor,  the  presence  or  absence  of  a  contact  agent  and 
the  nature  of  the  contact  agent,  as  weir  as  the  temperature.  No 
definite  kindling  temperature  can  be  given,  unless  the  other  con- 
ditions are  specified  also.  Kindling  conditions  involve  several 
variables,  of  which  the  temperature  is  only  one. 

Oxidation.  —  The  slower  union  with  oxygen  which  occurs  in 
rusting  is  called  oxidation.  We  shall  see  later,  however,  [that  it 
has  been  found  convenient  to  stretch  this  term  so  as  to  cover  com- 
binations of  other  elements  than  oxygen,  and  even  to  include 
actions  not  involving  combination.  At  this  point  we  can  discuss 
only  oxidation  by  oxygen. 

This  process  of  slow  oxidation  by  oxygen,  although  less  con- 
spicuous than  combustion,  is  really  of  greater  interest.  Thus  the 
decay  of  wood  is  simply  a  process  of  oxidation  whereby  the  same 
products  are  formed  as  by  the  more  rapid  ordinary  combustion. 
Sewage  is  mixed  with  large  volumes  of  river  water,  the  object  being, 
not  simply  to  dilute  the  sewage,  but  to  mix  it  with  water  containing 
oxygen  in  solution.  This  has  an  oxidizing  power  like  that  of  oxy- 
gen gas  and,  through  the  agency  of  bacteria,  quickly  renders  dis- 
solved organic  matters  innocuous  by  converting  them  for  the 
most  part  into  carbon  dioxide  and  water.  Thus,  a  few  miles 
further  down  the  stream,  the  water  becomes  as  suitable  for  drink- 
ing as  it  was  before  the  sewage  entered.  In  our  own  bodies  we 
have  likewise  a  familiar  illustration  of  slow  oxidation.  Avoiding 
details,  it  is  sufficient  to  say  that  the  oxygen,  from  the  air  taken 
into  the  lungs,  combines  with  the  haemoglobin  in  the  red  blood- 
corpuscles.  In  this  form  of  loose  combination,  it  is  carried  by  the 
blood  throughout  our  tissues  and  there  oxidizes  the  foodstuffs 
which  have  been  absorbed  during  digestion.  The  material  prod- 
ucts are  carbon  dioxide  and  water,  of  which  the  former  is  carried 
back  to  the  lungs  by  the  blood,  and  finally  reaches  the  ah-  during 


OXYGEN  37 

exhalation.     The  important  product,  however,  is  the  heat,  given 
out  by  the  oxidation,  which  keeps  the  body  warm. 

The  opposite  of  oxidation,  the  removal  of  oxygen,  is  spoken  of 
in  chemistry  as  reduction.  But  this  term,  also,  has  been  stretched 
to  cover  other  kinds  of  chemical  change. 

Spontaneous  Combustion.  —  Sometimes  a  mere  slow  oxi- 
dation develops  into  a  combustion,  which  is  then  known  as  spon- 
taneous combustion.  To  understand  this,  we  must  note  the  fact 
that  a  given  weight  of  material,  say,  iron,  in  combining  with  oxy- 
gen to  form  a  given  oxide,  will  liberate  the  same  total  amount  of 
heat  whether  the  union  proceeds  rapidly  or  slowly.  If  the  action 
proceeds  slowly,  and  the  material  being  oxidized  is  freely  exposed 
to  the  air,  the  latter  will  become  heated  and  will  carry  off  the  heat 
as  fast  as  it  is  produced.  Thus,  no  particular  rise  in  temperature 
will  occur.  If,  however,  the  material  is  a  poor  conductor  of  heat, 
like  hay  or  rags,  and  there  is  sufficient  air  for  oxidation,  but  not 
enough  to  carry  off  the  heated  air,  the  heat  may  accumulate  and  a 
temperature  sufficient  to  start  combustion  may  be  reached.  Such 
a  situation  sometimes  arises  in  hay-stacks.  It  occurs  also  when 
rags,  saturated  with  oils  used  in  making  paints  (linseed  oil  and 
turpentine)  are  left  in  a  heap.  These  oils,  in  "  drying,"  combine 
with  oxygen  from  the  air  and  turn  into  a  tough,  resinous  material. 
The  rags,  being  poor  conductors  of  heat,  finally  become  hot 
enough  to  burst  into  flame,  and  serious  conflagrations  often  owe 
their  origin  to  causes  such  as  this.  Oily  rags  should  always  be 
disposed  of  by  burning,  or  should  at  least  be  placed  in  a  closed  can 
of  metal.  Fires  in  coal  bunkers  of  ships  arise  from  the  same  cause 
—  slow  oxidation,  with  accumulation  of  the  resulting  heat.  That 
coal  does  undergo  slow  oxidation,  especially  when  freshly  mined, 
is  shown  by  the  fact  that  such  coal,  if  left  exposed  to  the  air  for. 
months,  may  lose  10  per  cent  or  more  of  its  heating  power. 

Uses  of  Oxygen.  —  A  number  of  the  practical  applications 
of  oxygen  have  already  been  mentioned.  For  example,  in  the 
foregoing  section  we  have  referred  to  its  use  in  breathing,  its  role 
in  decay,  which  is  a  beneficent  process  because  it  removes  much 
useless  matter  which  might  otherwise  cause  disease,  and  its  value 
in  the  disposal  of  sewage.  Power  and  heat  for  commercial  pur- 


38  COLLEGE    CHEMISTRY 

poses  are  almost  all  obtained  by  the  burning  of  coal,  in  which  oxy- 
gen from  the  air  plays  a  large  part.  If  we  had  to  purchase  the 
oxygen  as  well  as  the  coal,  we  should  require  at  least  three  tons  of 
oxygen  for  every  ton  of  coal. 

Oxygen  in  cylinders  and  oxygen  generators  are  used  to  restore 
the  supply  in  the  atmosphere  of  submarine  boats,  as  well  as  for  the 
purposes  already  mentioned  (p.  27). 

Substances  Indifferent  to  Oxygen.  —  Finally,  since  the 
atmosphere  contains  so  large  a  proportion  of  oxygen,  substances 
which  do  not  oxidize  and,  when  heated,  do  not  burn,  have  many 
uses.  Gold,  silver,  and  platinum  are  of  this  kind  (p.  32),  and  are 
used  for  ornaments.  The  last  is  used  for  crucibles  in  which  bodies 
are  heated  in  the  laboratory.  Although  iron  burns  in  pure  oxygen, 
it  does  not  oxidize  rapidly  in  the  air  even  when  heated,  and  so  is 
used  for  making  vessels  for  cooking  and  in  constructing  fireproof 
buildings. 

Compounds,  already  fully  oxidized,  are  naturally  not  com- 
bustible. Of  this  nature  are  sandstone,  granite,  brick,  porcelain, 
glass,  and  water.  All  these  are,  therefore,  fireproof.  Moreover, 
these  substances  do  not  give  off  oxygen  when  heated  (water  de- 
composes slightly).  Glass  and  porcelain  thus  neither  lose  nor 
gain  in  weight  when  heated,  and  are  suitable  materials  for  labora- 
tory apparatus. 

Activity  and  Stability.  —  A  substance  which  enters  into 
combination  vigorously,  as  does  oxygen,  is  called  chemically 
active.  Nitrogen,  on  the  other  hand,  is  relatively  inactive.  An 
active  element,  since  it  combines  eagerly,  naturally  holds  tena- 
ciously to  the  matter  with  which  it  has  combined.  An  active  ele- 
ment implies,  therefore,  also  one  which  is  in  general  difficult  to 
liberate  from  combination.  Its  compounds  are  in  general  rel- 
atively stable.  Thus,  many  oxides,  and  the  natural  compounds 
just  mentioned  (sandstone,  granite,  brick  and  porcelain,  the  last 
two  made  from  clay),  do  not  lose  oxygen  even  at  a  white  heat  and 
are  very  stable. 

Exercises.  —  1.  What  percentage  by  weight  of  free  oxygen  is 
obtained  by  heating:  (a)  mercuric  oxide,  (6)  potassium  nitrate, 


OXYGEN  39 

(c)  potassium  chlorate?    At  $1.50  (7/8),  $0.15  (8d),  and  $0.15  (8d) 
per  kilogram,  respectively,  which  is  the  cheapest  source  of  oxygen? 

2.  Using  the  data  on  pp.  30-31,  calculate  the  weight  of  oxygen 
dissolved  by  1000  c.c.  (=  1000  g.)  of  water  at  20°. 

3.  Why  does  a  forced  draft  make  a  fire  burn  more  rapidly? 

4.  Why  does  a  naked  flame  sometimes  cause  an  explosion  in  a 
mine,  when  the  air  of  the  mine  is  filled  with  coal  dust? 

5.  The  substances,  like  phosphorus  and  sulphur,  which  burn 
rapidly  in  ordinary  oxygen,  combine  very,  very  slowly  with  oxygen 
which  has  been  freed  from  moisture  by  careful  drying.    How  is 
this  effect  of  water  to  be  classified? 

6.  Air  is  20  per  cent  oxygen.    Why  does  iron  burn  brilliantly 
in  pure  oxygen,  but  not  in  air? 


CHAPTER  IV 

ATOMIC  WEIGHTS,    SYMBOLS,    FORMULA,   AND 
EQUATIONS 

WE  have  repeatedly  called  attention  to  the  quantities  of  the 
substances  taking  part  in  chemical  changes,  and  particularly  to 
the  constant  relation  between  the  weights  of  each  element  in  a 
given  substance  (pp.  17-18).  This  matter  is  of  great  importance  in 
chemistry.  If  a  cargo  of  copper  ore  is  to  be  purchased,  we  do  not 
wish  to  pay  for  the  rock  that  all  specimens  of  the  ore  contain  in 
larger  or  smaller  proportion.  So  we  secure  a  fair  sample  of  the  ore 
and  have  an  analysis  made  by  a  chemist.  The  analysis,  in  this 
case,  is  a  measurement  of  the  proportion  of  the  valuable  metal  in 
the  sample.  The  price  will  then  depend  largely  upon  the  propor- 
tion of  the  copper  per  ton  of  ore.  The  making  of  analyses  —  that 
is,  chemical  measurements  —  plays  a  very  large  part  in  all  in- 
dustries which  involve  the  consumption  or  manufacture  of  ma- 
terials. Quantitative  measurements,  aside  from  their  theoretical 
interest,  are  therefore  of  the  greatest  practical  importance.  Hence 
we  must  now  discuss  them  once  more. 

The  Compositions  of  Substances.  —  Our  present  purpose 
is  to  compare  the  proportions  by  weight  of  the  elements  composing 
several  compounds,  in  order  to  see  whether  the  numbers  are  really 
as  irregular  as,  in  the  examples  we  have  heretofore  given,  they 
have  appeared  to  be,  or  whether  there  is  any  way  of  relating  and 
simplifying  the  numbers. 

In  order  to  have  a  fair  sample  of  these  proportions,  we  shall  in- 
clude the  compositions  of  a  few  substances  for  which  the  data 
have  not  yet  been  given.  Potassium  hydroxide  (p.  35)  has  the 
composition:  potassium  (a  metal)  39.1,  oxygen  16,  hydrogen 
1.008,  in  a  total  of  56.108  parts.  Water  (oxide  of  hydrogen)  con- 
tains: oxygen  16  and  hydrogen  2.016  parts  by  weight.  When 
iron  burns  in  chlorine,  which  is  a  yellow  gas,  it  gives  ferric  chloride 

40 


ATOMIC   WEIGHTS,    SYMBOLS,    FORMULA,   AND   EQUATIONS     41 


with  the  proportions:  iron  55.84,  chlorine  106.38.  When  ferric 
chloride  is  heated  in  a  stream  of  hydrogen  gas,  a  part  of  the 
chlorine  is  removed,  and  ferrous  chloride  remains:  iron  55.84, 
chlorine  70.92. 

To  make  the  comparison  easy,  we  have  limited  the  number  of 
substances  to  five  of  those  previously  discussed,  together  with  the 
four  just  mentioned,  and  have  also  arranged  the  proportions  in 
the  form  of  a  table. 

PROPORTIONS  BY  WEIGHT  OF  THE  ELEMENTS  IN  CERTAIN 

COMPOUNDS 


Name  of  Compound. 

Iron. 

Oxy- 
gen. 

Sul- 
phur. 

Potas- 
sium. 

Chlo- 
rine. 

Hydro- 
gen. 

>>    Ferric  oxide  (p.  9)  
«>  Ferrous  sulphide  (p.  14)   .... 
1  ^Potassium  chlorate  (p.  27)  .    .    . 

"'^nlnhiir  HifiYiHp  (n    32^ 

111.68 
55.84 

48 

48 

32 

32^06 
32  06 

39!  i 

35  '.46 

.... 

Iron  oxide  (magnetic)  (p.  32)  .    . 
Potassium  hydroxide     
Water                

167.52 

64 
16 
16 

39  '.1 

i!66s 

2.016 

(    Ferric  chloride    

55.84 

106.38 

Ferrous  chloride 

55.84 

70.92 

Atomic  weights  

55.84 

16 

32.06 

39.1 

35.46 

1.008 

Study  of  the  Foregoing  Table.  —  When  we  first  examine 
the  numbers  in  the  horizontal  lines  of  the  table,  we  observe  that 
the  numbers,  with  the  exception  of  those  for  oxygen,  all  involve 
decimal  fractions.  From  this  we  infer  that  whole  numbers  must 
have  been  chosen  intentionally  for  oxygen.  This  is,  in  fact,  the 
case.  When  we  next  look  down  the  oxygen  column,  we  observe 
that  48  =  3  X  16  and  32  =  2  X  16  and  64  =  4  X  16.  All  the 
oxygen  weights  are  multiples  of  16  by  some  integral  (whole)  num- 
ber. In  the  hydrogen  column,  the  same  regularity  appears,  for 
2.016  =  2  X  1.008.  Following  up  this  idea,  we  find  in  the  iron 
column,  55.84  occurring  thrice,  and  discover  that  111.68  =  2  X 
55.84,  and  that  167.52  =  3  X  5£.84.  Similarly,  in  the  chlorine 
column,  the  numbers  are  multiples  of  35.46  by  unity  or  some  other 
integer.  Thus,  the  proportion  of  each  element,  in  various  com- 
pounds, can  be  represented  by  a  fundamental  number  —  a  sort  of 
unit  quantity  —  multiplied  when  necessary  by  the  proper  integer. 


42  COLLEGE    CHEMISTRY 

Now  this  rule  is  not  confined  to  these  nine  compounds,  involving 
only  six  different  elements.  If  we  provided  a  column  for  every 
known  element  (about  eighty  would  be  needed),  and  entered  the 
composition  of  every  known  compound,  we  should  find  the  same 
rule  to  hold.  This  rule  can  be  stated  as  follows: 

Law  of  Combining  Weights.  —  In  every  compound  sub- 
stance, the  proportion  by  weight  of  each  element  may  be  expressed 
by  a  fixed  number,  a  different  one  for  each  element,  or  by  a  mul- 
tiple of  this  number  by  some  integer  (whole  number). 

Since  the  proportion  by  weight  in  which  two  (or  more)  elements 
combine  is  a  chemical  property,  this  is  a  chemical  law.  Clearly,  it 
does  not  apply  to  mixtures,  for  any  irregular  proportion  could  be 
used  in  the  physical  process  of  mixing. 

Explanation  of  this  Law,  Atoms  and  Atomic  Weights.  — 

To  explain  this  law  it  was  necessary  to  use  the  third  kind  of  ex- 
planation (p.  23),  namely  the  making  of  an  hypothesis.  The 
details  of  how  two  substances  combine  cannot  be  seen,  so  chemists 
had  to  imagine  some  details  which  would  account  for  the  possession 
of  an  individual  unit  weight  by  each  element.  If  oxygen  is  com- 
posed of  minute,  invisible  particles,  which  are  all  alike  in  weight, 
and  hydrogen  and  potassium  are  of  the  same  nature,  except  that 
the  weight  of  the  particle  of  each  kind  of  element  is  different,  we  have 
the  basis  of  an  explanation.  We  have  to  suppose,  further,  that, 
when  elements  combine,  the  particles  adhere  in  pairs  or  groups,  as 
wholes,  and  are  never  broken.  In  this  way  the  particle  of  each 
variety  of  elementary  matter  will  have  a  definite,  unchangeable 
weight,  which  will  be  one  of  its  fixed  properties.  If  the  relative 
weights  of  the  particles  of  oxygen,  potassium,  and  hydrogen  are  in 
the  proportion  of  the  combining  numbers  in  the  table,  namely 
16  :  39.1  :  1.008,  the  whole  situation  becomes  clear.  Chemical 
union  must  consist,  in  detail,  in  the  union  of  the  particles  of  the 
elements  to  form  the  particles  of  the  compound.  For  each  particle 
of  potassium  hydroxide,  one  particle  each  of  the  three  elements 
is  required. 

For  each  particle  of  water,  where  the  proportion  of  oxygen  to 
hydrogen  is  16  :  2.016,  evidently  one  particle  of  oxygen  and  two 
particles  of  hydrogen  are  necessary.  Varying,  intermediate  pro- 


ATOMIC   WEIGHTS,    SYMBOLS,    FORMULAE,   AND   EQUATIONS     43 

portions  are  impossible,  because  the  particles  of  the  elements  are 
permanent,  are  never  broken,  and  combine  as  wholes,  and  in  a 
uniform  way  through  the  mass.  The  only  possible  variation 
would  be  to  take  different  relative  numbers  of  the  particles  —  for 
example,  two  of  oxygen  to  two  of  hydrogen  (2  X  16  :  2  X  1.008). 
But  this  product  would  have  a  different  composition  from  water, 
and  would  not  be  water.  This  compound,  with  the  double  pro- 
portion of  oxygen,  is  indeed  known  (it  is  hydrogen  peroxide),  and 
is  the  only  other  known  compound  of  these  two  elements. 

This  theory  fully  explains  why  the  combining  proportions  of 
each  element,  in  different  compounds,  can  always  be  expressed 
by  a  fixed,  unit  number  (which  represents  the  weight  of  the 
ultimate  particle  of  that  element),  multiplied,  when  necessary, 
by  a  whole  number  (representing  the  number  of  particles  of 
the  element  required  to  form  a  particle  of  the  compound  in 
question) . 

This  explanation  was  first  offered  by  Dalton,  a  schoolmaster 
of  Manchester  in  1802.  Borrowing  an  idea  from  the  speculations 
of  the  Greek  philosophers,  he  called  the  particles  of  elements 
atoms  (Gk.,  not  cut,  or  not  divided).  The  atoms  of  any  one  element 
are  all  alike  in  weight,  as  well  as  in  other  properties,  but  the  atoms 
of  different  elements  differ  in  weight. 

The  particles  made  by  uniting  two  or  more  atoms,  as  in  forming 
a  particle  of  a  compound,  are  called  molecules  (Gk.,  a  little 
mass) . 

A  chemical  combination  of  two  simple  substances  consists,  then, 
in  an  elaborate  re-grouping  of  the  atoms  of  both  elements  so  that 
molecules  of  the  compound  are  formed.  Definite  proportions  by 
weight  are  required,  in  order  that  the  atoms  of  each  element  may 
be  available  in  the  correct  proportion,  1  atom  :  1  atom,  or  1  :  2, 
or  2  :  3,  or  in  some  similar,  usually  simple  ratio. 

The  result  was  called  the  atomic  theory.  For  long  it  remained 
an  hypothesis,  or  sort  of  guess.  Recently,  however,  we  have 
obtained  independent  proof  that  molecules  and  atoms  are  real 
(see  Radioactivity),  for  we  can  now  count  and  measure  the  weight 
of  individual  molecules,  and  we  even  know  something  of  the  inside 
structure  of  atoms. 

The  fundamental  numbers,  one  for  each  element,  being  the 
relative  weights  of  the  atoms,  are  called  atomic  weights. 


44  COLLEGE    CHEMISTRY 

Symbols  and  Formulae.  —  One  self-evident  use  for  the 
atomic  weights  is  in  stating  the  compositions  of  compounds.  To 
make  the  statement  as  simple  as  possible,  symbols,  first  used  by 
Berzelius,  represent  the  atomic  weight  of  each  element.  Thus,  H 
stands  for  1.008  parts,  or  1  atom,  of  hydrogen,  and  0  for  16  parts, 
or  1  atom,  of  oxygen.  When  several  elements  have  the  same 
initial  letters,  another  letter  is  added:  C  for  one  atomic  weight  of 
carbon,  Ca  for  one  atomic  weight  of  calcium,  Cl  for  35.46  parts  of 
chlorine.  When  the  names  of  the.  elements  are  not  alike  in  all 
languages,  the  symbol  is  frequently  based  on  the  Latin  name,  as 
Fe  for  iron  (ferrum)  and  Pb  for  lead  (plumbum),  or  on  the  German, 
as  K  for  potassium  (kalium).  The  symbols  are  international.  A 
list  of  the  elements,  with  their  symbols  and  atomic  weights,  is 
printed  inside  the  back  cover  of  this  book. 

The  composition  of  any  compound  can  thus  be  stated  by  setting 
down  the  necessary  symbols,  together  with  the  whole  numbers,  if 
any,  by  which  the  atomic  weights  are  multiplied.  The  result  is  a 
formula.  For  example,  ferric  oxide  contains  iron  111.68  and  oxy- 
gen 48  parts  (p.  41).  This  is  equivalent  to  iron  2  X  55.84  and 
oxygen  3  X  16.  This  again  is  equivalent,  in  symbols,  to  2  X  Fe 
and  30.  The  formula  is  written  F^Os.  Ferrous  sulphide  is  a 
simpler  case:  iron  55.84  and  sulphur  32.06,  or,  in  symbols,  FeS. 
The  reader  should  now  examine  the  whole  table  on  p.  41,  and 
work  out  the  formula  of  each  compound  and  write  it  in  the 
margin. 

Equations.  —  It  is  now  possible  to  abbreviate  the  condensed 
statements  we  have  been  using  to  represent  the  substances  and 
their  quantities  in  chemical  reactions.  Thus,  the  three  statements 
on  p.  32,  when  translated  into  symbols,  are  as  follows: 

S  +  20  -»  S02. 

2P  +  50  -»  P205    (P  =  31.04). 
C  +  20->C02     (C  =  12). 

When  no  coefficient  appears  before  or  after  a  symbol,  1  is  to  be 
understood. 

Much  practice  is  required  to  enable  one  to  make  and  under- 
stand equations.  The  reader  should  therefore  at  once  turn  back 
to  the  statements  on  pp.  9,  14,  17,  27,  and  28,  obtain  the  necessary 


ATOMIC   WEIGHTS,    SYMBOLS,    FORMULA,    AND   EQUATIONS     45 

atomic  weights  and  symbols  from  the  table  at  the  end  of  the  book, 
and  construct  the  equation  in  each  case. 

The  term  "  equation"  refers  to  the  fact  that  the  total  weight 
of  matter  on  both  sides  is  always  the  same.  In  other  respects, 
such  as  in  the  nature  of  the  substances,  the  two  sides  are  entirely 
different. 

Derivation  of  Formulae  from  Experimental  Data.  —  In 

the  condensed  statements  referred  to  (by  page)  in  the  foregoing 
section,  the  numbers  given  were  already  multiples  of  the  atomic 
weights,  and  the  formulae  were  therefore  easy  to  make.  It  re- 
mains to  show  how  the  formula  may  be  constructed  from  the 
weights  obtained  in  an  experiment. 

In  the  quantitative  experiment  on  the  composition  of  cupric 
oxide  (p.  34),  the  proportion  found  was:  copper  85,  oxygen  21. 
In  the  formula,  the  same  proportion  is  to  be  expressed  by  means 
of  multiples  of  the  atomic  weights.  If  we  divide  each  of  these 
numbers  by  the  corresponding  atomic  weight,  the  quotient  will 
be  the  number  by  which  the  atomic  weight  must  be  multiplied. 
The  atomic  weights  are  Cu  =  63.57,  0  =  16.  85  -f-  63.57  =  1.3, 
and  21  -r-  16  =  1.3.  The  proportion  of  copper  to  oxygen  in  the 

,    85  63.57  X  1.3 

compound,  ^  ,  now  becomes     1ft      1  Q    • 

2i\.  1O  X  l.o 

But  this  proportion  must  be  expressed  in  multiples  of  the  atomic 
weights  by  whole  numbers.     Dividing  above  and  below  by  1.3,  we 
.  63.57  X  1 

get  -16^1- 

Now  the  symbols  stand  for  the  atomic  weights.  Substituting 
the  symbols,  the  proportion  becomes  •„  .  The  formula  is, 

therefore,  CuO. 

Applying  the  same  process  to  the  case  of  sulphur  dioxide  (p.  35)  : 


Sulphur  _  32.06  _  32.06  X  1      S  X  1 
Oxygen  =     32  16  X  2         O  X  2' 


If  the  composition  of  the  substance  has  been  stated  in  percent- 
ages, the  same  device  is  used.  Thus,  the  case  of  sodium  sulphate 
works  out  as  follows: 


46 


COLLEGE    CHEMISTRY 


Elements. 

Percentages. 

At.  Wt.     Quotient 

+  ' 

Formula. 

Sodium     
Sulphur    

32.43 
22.55 

23    X    1.41 
32  X  0.705 

0.705 
0.705 

NaX  2 

s 

Oxygen     

45.02 

16  X  2.814 

0.705 

O  X  4 

The  formula  is,  therefore, 

It  is  obvious  that,  after  we  have  found  out  what  elements  com- 
pose a  given  compound,  we  are  still  unable  to  write  its  formula. 
We  may  not  simply  set  the  symbols  down,  side  by  side.  A  meas- 
urement must  be  made,  in  order  that  we  may  find  out  the  factors 
by  which  the  atomic  weights  are  to  be  multiplied. 

Answers  to  Some  Questions.  —  Why  was  a  whole  number 
assigned  to  oxygen?  Oxygen  was  chosen  as  the  basis  of  the  system 
because  the  exact  determinations  of  the  combining  weights  of  most 
of  the  elements  have  actually  been  made  by  direct  union  with 
oxygen,  or  with  the  help  of  but  one  intermediate  step.  If  the 
question  had  been  one  of  mathematics,  hydrogen,  the  element 
with  the  lowest  combining  proportions,  would  have  furnished  the 
basis  and  unit  of  the  scale.  But  the  question  was  the  practical 
one  of  getting  the  most  accurate  measurements  for  the  relative 
magnitudes  of  the  numbers.  Hydrogen  combines  with  only  a  few 
of  the  elements,  and  the  proportion  of  hydrogen  is  usually  so  small 
that  the  weights  of  this  element  cannot  be  measured  so  accurately 
as  can  the  much  larger  weights  of  oxygen  and  of  the  other  elements. 
So  oxygen  was  selected  as  the  basal  element. 

Why  was  16  assigned  to  oxygen,  rather  than  1  or  100,  or  some 
other  whole  number?  The  number  16  was  chosen  in  order  that 
the  advantage  of  having  a  mathematical  unit,  or  something  close 
to  it,  in  the  scale,  might  be  retained  also.  With  this  value,  hydro- 
gen became  1.008.  A  whole  number  smaller  than  16  would  make 
the  atomic  weight  of  hydrogen  less  than  unity.  With  H  =  1, 
the  value  for  oxygen  becomes  about  15.9,  and  the  values  for  all  the 
elements  are  changed  in  proportion.  The  result  of  such  a  change 
would  be  that  the  values  for  the  common  elements  would  not  be  so 
close  to  whole  numbers  as  they  at  present  are  (e.g.,  C  =  12.00, 
N  =  14.01,  Na  (sodium)  =  23.00,  K  =  39.1,  P  =  31.04).  With 


ATOMIC   WEIGHTS,    SYMBOLS,    FORMULA,    AND   EQUATIONS     47 

O  =  16,  it  is  possible,  and  of  course  more  convenient,  in  many 
cases  to  use  the  nearest  whole  number  in  ordinary  calculations. 

The  answers  to  the  two  foregoing  questions  show  why  the  scale 
of  the  numbers  was  fixed  as  it  is.  Of  course,  multiplying  or 
dividing  all  the  atomic  weights  by  any  number,  whole  or  fractional, 
would  not  affect  their  scientific  accuracy.  The  choice  of  scale  is 
merely  a  matter  of  convenience. 

In  physics  there  is  one  unit  of  weight,  the  gram,  for  all  kinds  of 
matter.  Is  it  the  case  that  in  chemistry  a  different  unit  of  weight 
is  employed  for  each  element?  This  is  the  exact  situation,  and 
it  is  one  peculiar  to  chemistry.  It  does  not  represent  an  arbitrary 
decision  of  the  chemist,  however.  It  is  due  to  the  fact  that  the 
atoms  of  any  one  element  have  the  same  weight,  but  that  the 
atoms  of  different  elements  have  different  weights.  The  atom  of 
uranium  is  238  times  as  heavy  as  that  of  hydrogen,  and  its  com- 
bining proportions,  therefore,  are  in  general  greater  in  the  same 
ratio,  while  the  atoms  of  the  other  elements  have  weights  falling 
between  these  limits. 

There  is  still  one  question  to  be  asked.  Why  take  16  for  oxy- 
gen rather  than  8  or  32?  In  other  words,  may  we  not  multiply  or 
divide  any  one  (or  more)  of  the  individual  atomic  weights  by  a 
whole  number?  The  answer  is  that,  thus  far,  we  have  not  met 
with  any  reason  for  not  doing  so.  With  O  =  8,  and  H  still  1.008, 
the  composition  of  water  would  be  represented  by  the  formula  HO 
instead  of  H20  (where  0  =  16).  In  a  later  chapter  (Chap.  VIII), 
however,  we  shall  see  that  the  individual  numbers  actually  chosen 
meet  certain  other  conditions,  in  addition  to  those  already  men- 
tioned, and  are  on  that  account  preferable  to  any  other  set. 

Law  of  Multiple  Proportions.  —  We  have  already  met  with 
several  instances  in  which  two  elements  combine  in  more  than 
one  proportion  by  weight,  and  form  therefore  more  than  one  com- 
pound. Thus  two  oxides  of  iron  and  two  chlorides  of  iron  have 
been  mentioned  (p.  41),  and  two  oxides  of  hydrogen,  water  and 
hydrogen  peroxide,  are  known  (p.  43).  This  general  fact  was  dis- 
covered before  the  law  of  combining  weights  (p.  42)  had  been  for- 
mulated, and  is  a  particular  case  of  this  law.  It  was  discovered 
by  Dalton  (1804)  and  was  embodied  by  him  in  a  statement  known 
as  the  law  of  multiple  proportions,  which  ran  somewhat  as  follows : 


48  COLLEGE    CHEMISTRY 

If  two  elements  unite  in  more  than  one  proportion,  forming  two  or 
more  compounds,  the  quantities  of  one  of  the  elements,  which  in  the 
different  compounds  are  united  with  identical  amounts  of  the  other* 
stand  to  one  another  in  the  ratio  of  integral  numbers,  which  a$e 
usually  small. 

The  two  chlorides  of  iron  illustrate  the  law.  Ferric  chloride 
contains  iron  55.84  and  chlorine  106.38,  and  ferrous  chloride  iron 
55.84  and  chlorine  70.92.  Thus  the  quantities  of  chlorine  united 
with  identical  amounts  of  iron  (namely,  55.84  parts)  stand  in  the 
ratio  106.38  :  70.92,  or  3  :  2. 

Exercises.  —  1.  From  the  data  on  p.  9  and  the  atomic  weights, 
calculate  the  formula  of  lead  oxide.  Construct  also  the  equations 
for  the  decomposition  of  potassium  chlorate  (p.  27),  and  for  the 
action  of  water  in  sodium  peroxide  (p.  28). 

2.  When  1  g.  of  sodium  burns  in  oxygen,  it  produces  1.7  g.  of 
the  oxide.     What  is  the  formula  of  the  latter  and  the  equation? 

3.  If  26  g.  of  mercurous  oxide  are  required  to  give,  by  heating, 
1  g.  of  oxygen,  what  is  the  formula  of  the  substance? 

4.  What  are  the  formulae  of  the  substances  possessing  the  fol- 
lowing percentage  compositions? 

I  II  III 

Magnesium,  25.57  Sodium,   32.43  Potassium,  26.585 

Chlorine,        74.43  Sulphur,  22.55  Chromium,  35.390 

Oxygen,    45.02  Oxygen,       38.025 

5.  What  are  the  percentage  compositions  of  substances  possess- 
ing the  following  formulae:   Mn304,   KBr,    FeSO4? 

6.  Compare  the  formula  of  mercurous  oxide,  found  in  3,  with 
that  of  mercuric  oxide,  and  show  how  the  compounds  illustrate 
the  law  of  multiple  proportions  (p.  48). 

7.  If  the  atomic  weight  of  potassium  were  13.03,  and  the  other 
atomic  weights  were  unchanged,  what  would  be  the  formulae  of 
(a)  potassium  hydroxide,  and  (6)  potassium  chlorate? 


CHAPTER  V 
HYDROGEN 

HAVING  learned  something  of  the  nature  of  the  atmosphere, 
and  particularly  of  oxygen,  its  most  active  component,  we  turn 
now  to  water,  a  substance  as  closely  connected  with  our  daily  life 
as  is  air.  We  find  that  it  is  a  compound  of  oxygen  and  hydrogen, 
and  the  latter  element,  therefore,  may  be  taken  up  next.  Hydro- 
gen is  of  interest  on  its  own  account  because  it  is  often  used  in 
filling  balloons,  and  nearly  half  the  bulk  of  ordinary  illuminating 
gas  is  free  hydrogen. 

History.  —  That  hydrogen  is  a  distinct  kind  of  gas  was  first 
established  by  Cavendish  (1766).  Somewhat  later  (1781),  he 
showed  that,  when  it  burned  in  the  air,  it  gave  a  vapor  which 
could  be  condensed  to  liquid  water.  Since  oxygen  was  then  known 
to  be  the  substance  with  which  combustibles  united,  this  proved 
that  water  was  a  compound  of  hydrogen  (Gk.,  water  producer)  and 
oxygen. 

Occurrence.  —  Free  hydrogen  is  found,  mixed  with  varying 
proportions  of  other  gases,  in  exhalations  from  volcanoes,  in 
pockets  found  in  certain  layers  of  the  rock-salt  deposits,  and  in 
some  meteorites.  The  air  contains  not  over  1  part  in  1,500,000. 
The  lines  of  hydrogen  are  prominent  in  the  spectra  of  the  sun  and 
of  most  stars. 

In  combination,  it  constitutes  about  11  per  cent  of  water.  It  is 
an  essential  constituent  of  all  acids.  It  is  contained  also,  in  com- 
bination with  carbon,  in  the  components  of  natural  gas,  petroleum, 
and  all  animal  and  vegetable  bodies. 

Preparation  by  the  Action  of  Metals  on  Cold  Water.  — 

To  liberate  hydrogen  from  water,  it  is  necessary  to  use  some  ele- 
ment with  which  the  oxygen  of  the  water  will  combine  even  more 
eagerly  than  with  hydrogen,  and  to  offer  this  element  in  exchange 
for  the  hydrogen. 

49 


50  COLLEGE    CHEMISTRY 

The  more  active  metals,  such  as  potassium  (K),  sodium  (Na), 
or  calcium  (Ca),  displace  hydrogen  rapidly  from  cold  water.  Po- 
tassium and  sodium  are  lighter  than  water,  and  float  on  the  sur- 
face. In  the  case  of  the  former,  so  much  heat  is  liberated  that  the 
hydrogen  catches  fire,  and  with  neither  metal  is  the  experiment 
safe  in  the  hands  of  a  novice.  Calcium  sinks  to  the  bottom,  and 
acts  rapidly,  but  not  violently,  so  that  the  gas  is  easily  collected 
(Fig.  22).  The  pieces  of  these  metals,  of  course,  act  upon  only  a 
small  part  of  the  water  in  the  vessel.  In  each  case 
the  metal  displaces  one-half  only  of  the  hydrogen  in 
that  part  of  the  water  upon  which  it  acts.  The 
other  products  are  the  hydroxides  of  potassium, 
sodium,  and  calcium,  respectively.  The  two  former 
dissolve,  leaving  a  clear  liquid  when  the  metal  is 
all  gone,  but  may  be  recovered  as  white  solids  by 
evaporation.  The  calcium  hydroxide  (slaked  lime) 
is  dissolved  only  in  part,  and  much  of  it  may  be  seen 
suspended  in  the  water  after  the  action. 

An  alloy  of  lead  with  sodium  (35  per  cent),  sold  under  the  name 
of  hydrone,  affords  a  convenient  substitute  for  sodium  in  the  fore- 
going actions. 

The  Making  of  Equations.  —  To  make  an  equation  we  must 
have  the  results  of  quantitative  measurements.  These  furnish 
us  with  the  composition  of  each  substance  concerned.  The  com- 
position, expressed  in  multiples  of  the  atomic  weights,  is  recorded 
in  the  formula  for  the  substance.  If  we  are  in  possession  of  the 
necessary  formulse,  we  can  write  the  equation. 

For  example,  the  composition  of  water  is:  hydrogen  2  X  1.008, 
oxygen  16.  In  symbols,  this  is  2H  and  0,  and  the  formula  is, 
therefore,  H20.  The  composition  of  potassium  hydroxide  is: 
potassium  39.1,  oxygen  16,  hydrogen  1.008,  and  the  formula,  there- 
fore, KOH.  In  calcium  hydroxide  the  proportions  are:  calcium 
40.07,  oxygen  2  X  16,hydrogen  2  X  1.008,  andthe  formula  Ca(OH)2. 

To  make  the  equation,  we  first  write  down  the  formulse  of  the  sub- 
stances used  and  produced: 

K  +  H20  ->  KOH  +  H. 
Na  +  H20  ->  NaOH  -f  H. 
Skekton:  Ca  +  H20  ->  Ca(OH)2  +  H. 


HYDROGEN 


51 


Next  we  must  balance  this  equation,  if  necessary.  That  is,  we 
must  adjust  it  so  that  there  are  equal  numbers  of  atomic  weights 
(or  atoms)  of  each  element  on  both  sides  of  the  equation.  This  is 
necessary  only  in  the  third  equation,  and  is  done  because,  accord- 
ing to  the  law  of  conservation  of  mass,  there  must  be  the  same 
quantity  of  each  element  after  the  reaction  as  there  was  before  it. 
On  examining  the  third  equation,  we  note  that  there  is  20,  in  the 
(OH) 2,  on  the  right  side  and  only  0  on  the  left.  We  therefore 
place  a  2  in  front  of  the  H2O,  for  we  cannot  get  the  additional  oxy- 
gen excepting  by  using  more  water: 
Balanced:  Ca  +  2H20  ->  Ca(OH)2  +  2H. 

The  number  of  atomic  weights  of  hydrogen  is  made  equal  by  using 
2H  on  the  right  side. 

The  coefficients  in  front  of  a  formula  multiply  the  whole  formula. 
Thus,  2H2O  is  equivalent  to  2(H20).  A  subscript  coefficient 
following  a  symbol,  however,  multiplies  that  symbol  only.  Thus 
H20  is  equivalent  to  (H)2O,  or  (2  X  H  +  0). 

Four  Steps  in  Making  an  Equation.  —  1.  Find  out,  by 
observation  and  experiment,  what  substances  are  used  and  what 
substances  are  produced. 

2.  Find  the  formula  of  each  substance  used  or  produced. 

3.  Set  the  formulae  down  as  a  skeleton  equation,  placing  the 
formulae  of  the  substances  used  on  the  left,  and  of  those  produced  on 
the  right. 

4.  Adjust,  or  balance  the  equation,  if  necessary. 
The  reader  must  practice  the 

making  of  equations,  until  he  can 
do  it  quickly.  The  text  contains 
many  equations,  but  more  usu- 
ally only  the  data  required  for 
making  them  (the  formulae  of  the 
substances)  are  given. 

Hydrogen  from  Metals  and 
Water  at  a  High  Tempera- 
ture.—  With  steam  at  a  red 
heat,  metals  like  iron,  zinc,  and  magnesium  interact  vigorously.  The 
steam,  generated  in  a  flask,  enters  at  one  end  of  the  tube  containing 
the  metal  (Fig.  23) ,  and  the  hydrogen  passes  off  at  the  other.  Since, 


FIG.  23. 


52  COLLEGE    CHEMISTRY 

at  a  red  heat,  all  hydroxides,  except  those  of  potassium  and  sodium, 
are  decomposed  into  an  oxide  of  the  metal  and  water,  as,  for 
example,  Mg(OH)2  — »  MgO  -f  H20,  the  oxides  are  formed  in  this 

case: 

Mg-{-H20-»Mg04- 2H. 

Iron  gives  the  magnetic  oxide  FcsC^. 

Making  Equations,  Again.  —  The  skeleton  equation  for  the 
action  of  iron  on  steam  is: 

Skeleton:  Fe  +  H20  ->  Fe304  +  H. 

We  are  not  permitted  to  alter  these  formulae  themselves,  but  we 
may  put  coefficients  in  front  of  any  of  them  to  make  the  number 
of  atomic  weights  alike  on  both  sides.  A  useful  rule  is  to  pick  out 
the  largest  formula  and  reason  back  from  that.  Here,  this  is 
FesO-j.  To  get  Fe3,  we  must  start  with  3Fe,  and  to  get  O4,  we  must 
start  with  4H20: 

Balanced:  3Fe  +  4H20  -»  Fe304  +  8H. 

Acids.  —  In  making  hydrogen  in  the  laboratory,  the  acids  are 
used  almost  exclusively.  The  common  acids  are  hydrochloric  acid 
(HC1,  Aq),  and  sulphuric  acid  (H2S04,  Aq).  The  usual  forms  are 
mixtures  containing  water,  the  variable  amount  of  the  latter  being 
indicated  by  the  symbol  Aq.*  The  former  is  a  solution  of  a  gas, 
hydrogen  chloride.  The  "pure  concentrated''  hydrochloric  acid 
used  in  laboratories  contains  nearly  as  much  of  the  gas  (39  per 
cent  by  weight)  as  the  water  can  dissolve.  The  " commercial" 
acid  contains  impurities  and  is  also  less  concentrated.  The  "con- 
centrated" sulphuric  acid  is  an  oily  liquid  containing  practically 
no  water.  The  "commercial"  sulphuric  acid  contains  6  to  7  per 
cent  of  water,  besides  impurities.  Acetic  acid  (HCO2CH3,  Aq)  is  a 
solution  of  a  liquid  in  water,  and  is  the  acid  found  in  vinegar. 

All  the  "dilute"  acids  contain  70  to  80  per  cent  of  water.  The 
water,  as  a  rule,  takes  no  part  in  the  chemical  changes  in  which  the 
acids  are  concerned,  and  is  therefore  omitted  from  the  equations. 

*  The  formula  H20  stands  for  a  fixed  proportion  of  water,  namely  18  parts. 
The  water  in  these  solutions  is  not  combined,  and  can  be  varied  in  amount,  so 
that  the  formula  H2O  may  not  logically  be  employed  here. 


HYDROGEN  53 

The  name  "acid"  is  restricted  to  one  class  of  substances  having 
certain  definite  characteristics.  Hydrogen  is  the  one  essential  con- 
stituent of  all  acids.  Their  aqueous  solutions  have  a  sour  taste  and 
change  the  color  of  litmus  from  blue  to  red.  When  free  from  water 
they  do  not  conduct  electricity.  When  dissolved  in  water  they 
conduct,  and  are  decomposed  by  the  electric  current.  In  aqueous 
solution,  also,  their  hydrogen  (or  one  unit  weight  of  it  in  the  case 
of  acetic  acid)  is  displaced  by  certain  metals. 

Radicals.  —  In  describing  the  chemical  behavior  of  acids,  we 
speak  of  the  hydrogen  as  the  positive  radical,  because  in  electrolysis 
(see  p.  55)  it  is  attracted  to  the  negative  pole,  and  of  the  material 
combined  with  the  hydrogen  as  the  negative  radical,  because  it  is 
attracted  to  the  positive  pole.  Thus  the  negative  radicals  in  the 
above  acids  are  Cl,  S04,  and  C02CH3,  respectively.  The  first  (Cl) 
is  a  simple  radical,  the  others  are  compound  radicals.  In  many 
interactions  the  compound  radicals  move  as  units  from  one  state 
of  combination  to  another. 

Preparation    by    Displacement  from    Diluted  Acids. — 

Every  one  of  the  metals  which  displace  hydrogen  from  water  will 
also  displace  it  from  dilute  acids.  The  acids  must  be  diluted  with 
water,  unless,  like  hydrochloric  acid,  they  are  already  dissolved  in 
water.  The  action  is  much  more  vigorous  than  that  on  water,  so 
that  the  most  active  metals  are  not  employed.  Metals  like  zinc, 
iron,  and  aluminium  serve  the  purpose.  The  metal  combines 
with  the  negative  radical,  and  so  liberates  the  hydrogen,  which 
escapes  in  bubbles.  Evaporation  of  the  clear  liquid,  when  the 
metal  has  all  disappeared,  gives  in  dry  form  the  compound  of  the 
metal  with  the  negative  radical.  Thus,  with  zinc  and  dilute 
sulphuric  acid,  zinc  sulphate  ZnS04  is  produced. 

Skeleton:  Zn  +  H2S04  -» ZnSO4  +  H. 

Balanced:  Zn  +  H2S04  -» ZnS04  +  2H. 

With  aluminium  and  hydrochloric  acid,  the  product  is  aluminium 

chloride  A1C13: 

Skeleton:  Al  +  HC1    ^A1C13  +  H. 

Balanced:  Al  +  3HC1  -»  A1C13  +  3H. 


54 


COLLEGE    CHEMISTRY 


The  water  undergoes  no  change  during  the  action,  although  its 
presence  is  essential.  It  is  simply  a  part  of  the  apparatus.  Any 
acid  may  be  used,  although  with  many  the  action  goes  on  very 
slowly. 

For  preparing  small  amounts  of  hydrogen,  the  apparatus  (Fig. 
24)  is  such  that  additional  acid  may  be  added  through  the  thistle-, 


0 


FIG.  24. 


Fia.  25a. 


FIG.  25b. 


or  safety  tube.     This  avoids  opening  the  flask  and  admitting  air. 

The  gas  may  be  caught  like  oxygen  over  water  or,  being  lighter 
than  air,  may  be  collected  by  downward  displace- 
ment of  the  latter  (Fig.  25a).  Heavy  gases  are 
collected  by  upward  displacement  of  air  (Fig. 
25b). 

With  a  Kipp's  apparatus  (Fig.  26)  the  gas 
may  be  made  on  a  large  scale  and  its  delivery 
can  be  regulated.  When  the  stream  of  gas  is 
shut  off  by  the  stopcock,  the  pressure  of  the 
gas,  as  it  continues  to  be  generated,  drives  the 
acid  away  from  the  metal  and  up  into  the 
globe  above,  so  that  the  action  ceases.  Yet 
the  action  is  ready  to  begin  again  the  moment 
any  portion  of  the  stored  gas  is  drawn  off  for  use. 
Silver,  gold,  and  platinum,  which  do  not 
combine  with  free  oxygen,  and  even  copper 
and  mercury,  which  do,  are  all  unable  to  lib- 
FIG.  26.  erate  hydrogen  and  to  form  oxides  when  heated 

in  steam.     When  treated  with  dilute  acids,  none  of  these  metals 

is  able  to  displace  and  liberate  the  hydrogen  (see  order  of  activity 

of  the  metals,  p.  59). 


HYDEOGEN 


55 


Contact  of  the  zinc  or  iron  with  an  inactive  metal,  like  platinum 
or  copper,  forms  an  electrical  couple  and  hastens  the  interaction. 
For  the  same  reason,  commercial  zinc,  which  contains  traces  of 
other  metals,  gives  a  steady  evolution  of  hydrogen,  while  extremely 
pure  zinc  is  almost  inactive. 

The  Third  Variety  of  Chemical  Change:  Displacement. 

—  The  reactions  used  in  liberating  hydrogen  illustrate  the  third 

of  the  four  common  forms  of  chemical  change.     Here  a  simple 

substance  (the  metal)  and  a  compound  (the  acid)  interact;  the 

compound  is  divided  into  its  radicals; 

and  the   simple   substance   combines 

with  one  radical  while  the  other  radical 

is  liberated.    The  interacting  element, 

here  the  metal,  is  said  to  displace  the 

other  element,  here  the  hydrogen,  from 

combination.   The  action  of  metals  on 

water  is  a  displacement  also. 

Preparation  of  Hydrogen  by 
Electrolysis.  —  If  we  dissolve  any 
acid  in  water,  and  immerse  the  wires 
from  a  battery  in  the  solution,  bubbles 
of  hydrogen  begin  to  appear  on  the 
negative  wire  (the  cathode)  and  rise  to 
the  surface.  All  the  other  constituents, 
whatever  they  may  be,  are  attracted 
to  the  positive  wire  (the  anode)  and, 
therefore,  do  not  interfere  with  the 
collection  of  pure  hydrogen.  An  appa- 
ratus devised  by  Hofmann  (Fig.  27) 
enables  us  to  secure  the  hydrogen, 
which  ascends  on  the  left  and  accumu- 
lates at  the  top  of  the  tube,  displacing  Fia-  27- 
the  solution.  When  hydrochloric  acid  is  used:  HC1  — » H  (neg.  wire) 
+  Cl  (pos.  wire),  the  chlorine,  a  soluble  gas,  remains  dissolved  in 
the  water  near  the  positive  pole.  When  sulphuric  acid  is  employed: 


H2S04  -»  2H  (neg.  wire)  +  S04  (pos.  wire). 


(1) 


56  COLLEGE   CHEMISTRY 

The  S04,  however,  acts  upon  the  water: 

S04-fH20-»H2S04  +  0.  (2) 

Thus,  the  sulphuric  acid  is  re-formed,  round  the  positive  wire, 
and  only  hydrogen  and  oxygen  are  finally  liberated. 

Decomposition  of  a  compound  by  the  use  of  electrical  energy  is 
called  electrolysis  (Gk.,  decomposition  by  electricity). 

The  Other  Ways  of  Preparing  Hydrogen.  —  For  special 
purposes,  hydrogen  may  be  made  by  boiling  an  aqueous  solution 
of  sodium  hydroxide  with  aluminium  turnings,  when  sodium  alumi- 
nate  is  formed:  Al  +  NaOH  +  H2O  ->  NaAlO2  +  3H;  also  by 
heating  powdered  zinc  and  dry  sodium  hydroxide,  the  product 
being  sodium  zincate:  Zn  +  2NaOH  ->  Na2Zn02  +  2H. 

Sources  of  the  Hydrogen  of  Commerce.  —  Zinc  is  too 
expensive  a  substance  to  use  in  the  preparation  of  hydrogen  in 
large  quantities  for  commercial  purposes.  We  realize  this  when 
we  note  that  33  parts  of  zinc  will  liberate  only  one  part  of  hydrogen, 
so  that  with  1  Ib.  of  zinc  we  obtain  only  one  half-ounce  of  the  gas. 
Different  sources  are  used  in  different  localities  and  countries. 

The  largest  supply  is  probably  obtained  as  a  by-product  in  the 
electrolysis  of  an  aqueous  solution  of  common  salt  (NaCl),  in 
connection  with  the  manufacture  of  caustic  soda  (sodium  hy- 
droxide, q.v.).  The  hydrogen  is  collected  and  compressed  in  steel 
cylinders. 

In  some  circumstances,  the  method  of  passing  steam  over  heated 
iron  is  used  (p.  51). 

Another  plan  is  to  liquefy  water-gas  (q.v.),  a  mixture  of  hydrogen 
and  carbon  monoxide.  The  hydrogen  evaporates  much  the  more 
readily  of  the  two,  and  can  thus  be  separated.  This,  and  still 
other  processes,  involve  substances  and  reactions  which  we  have 
not  yet  encountered  and  will  be  mentioned  at  the  appropriate  points. 

Physical  Properties.  —  Hydrogen  is  a  colorless,  tasteless, 
odorless  gas.  One  liter  weighs  only  0.08987  g.,  while  one  liter  of 
air  weighs  1.293  g.  Air  is  thus  14.5  times  heavier,  and  hydrogen 
can  be  poured  upwards  (Fig.  28)  and  is  used  for  filling  balloons. 
Hydrogen  was  first  liquefied  in  visible  amounts  by  Dewar  (1898). 


HYDROGEN 


57 


The  critical  temperature  is  -234°.  The  colorless  liquid  boils  at 
—  252.5°  and,  when  allowed  to  evaporate  rapidly  under  reduced 
pressure/  freezes  to  a  color- 
less solid  (m.-p.  -260°).  All 
other  gases,  except  helium, 
solidify  easily  when  led  into  a 
vessel  surrounded  by  liquid 
hydrogen. 

Hydrogen  is  even  less  sol- 
uble in  water  than  is  oxygen, 
1.8  volumes  of  the  gas  dissolve 
in  100  volumes  of  water  at  15°. 
Hydrogen  is  absorbed,  for  the 
most  part  in  a  purely  mechan- 
ical way,  by  many  metals. 
Heated  iron  will  take  up  19  FIG.  28. 

times  its  volume  of  hydrogen,  gold  takes  up  46  volumes,  platinum 
in  fine  powder  50  volumes,  palladium  502  volumes,  and  silver  none. 
The  maximum  absorbed  by  palladium  under  favorable  conditions 
is  873  volumes. 

Diffusion.  —  When  two  cylinders,  one  filled  with  hydrogen 
and  one  with  air,  are  placed  mouth  to  mouth  (Fig.  29),  so  that  the 
one  containing  hydrogen  is  uppermost,  since  the  air 
in  the  lower  cylinder  is  14.5  times  heavier  than  the 
hydrogen,  we  might  expect  the  gases  to  remain  in 
their  respective  cylinders.  The  air,  however,  makes 
its  way  into  the  hydrogen  above  it,  and  the  hydrogen 
penetrates  into  the  air  in  the  lower  cylinder  so  that, 
in  a  short  time,  the  gases  are  perfectly  mixed,  just  as 
if  gravity  did  not  exist.  The  same  phenomenon  is 
observed  when,  in  everyday  life,  a  bottle  of  scent  is 
opened.  The  vapor,  on  escaping,  begins  to  penetrate  in  all 
directions  through  the  room,  showing  its  presence  by  its  odor. 
The  material  of  gases  has  in  fact  an  independent  power  of  loco- 
motion. The  resulting  phenomenon  we  call  diffusion.  It  is 
constant  in  rate  for  each  gas  under  like  conditions,  and  hydrogen 
has  the  greatest  speed  of  diffusion  of  all  the  gases. 

The  different  rates  of  diffusion  of  different  gases  are  easily 


FIG.  29. 


58 


COLLEGE    CHEMISTRY 


shown  by  comparing  their  several  speeds  with  that  of  air,  when 
both  pass  through  a  wall  of  unglazed,  porous  porcelain. 

The  porous  cylinder  A  (Fig.  30)  contains  air  and  is  connected  by 
a  rubber  stopper  with  a  wide  tube  which  dips  beneath  the  surface 
of  the  water.  When  a  cylinder  H  containing  hydro- 
gen is  brought  over  it,  rapid  escape  of  gas  takes  place 
through  the  water,  showing  that  a  rise  in  pressure  has 
taken  place  inside  the  porous  vessel.  Before  the  cylin- 
der of  hydrogen  approached  the  porous  vessel,  the  air 
was  moving  both  outwards  and  inwards  through  the 
porcelain,  but,  being  the  same  air,  the  speed  of  motion 
was  equal  in  both  directions,  and  therefore  the  pres- 
sure inside  was  not  affected.  It  is  important  to  note 
that  there  was  at  no  time  rest,  there  was  simply  equal 
motion  in  both  directions.  When  the  hydrogen  at- 
mosphere surrounded  the  cylinder,  the  hydrogen  gas 
moved  more  rapidly  into  the  cylinder  than  the  air  in- 
side could  move  out,  and  hence  an  excess  of  pressure 
quickly  arose  in  the  interior. 

Exact  measurement  shows  that  the  lighter  a  gas  is 
in  bulk,  the  faster  its  parts  move  by  diffusion  in  any 
direction.  The  rate  is  inversely  proportional  to  the  square  root  of 
the  density  of  the  gas.  Thus,  for  hydrogen  and  air  it  is  in  the  ratio 
Vl.293  :  V00897,  or  3.8  :  1. 


FIG.  30. 


Chemical  Properties.  —  Hydrogen, 
delivered  from  a  jet,  burns  in  air  or  pure 
oxygen.  A  cold  vessel  held  over  the 
almost  invisible  blue  flame  condenses  to 
droplets  of  water  the  steam  that  is  pro- 
duced (Fig.  31).  When  hydrogen  and 
oxygen  are  mingled  in  a  suitable  burner 
(Fig.  32),  although  the  flame  gives  little 
light,  it  is  exceedingly  hot.  Platinum 


J 


II 


FIG.  31. 


FIG.  32. 

melts  in  it  easily  and  an  iron  wire  burns  brilliantly.  In  a  closed 
space  it  produces  a  temperature  of  over  2500°.  When  the  flame  is 
allowed  to  play  on  a  piece  of  quicklime,  the  latter  becomes  white- 
hot  at  the  spot  where  the  flame  meets  it.  This  result  is  called  a 
calcium  light  or  lime  light. 


HYDROGEN  59 

When  hydrogen  and  oxygen  are  mixed,  the  chemical  action  is 
very  slow  at  ordinary  temperatures,  no  perceptible  amount  of  union 
occurring  in  a  period  of  five  years.  If  the  mixture  is  sealed  up  and 
kept  at  300°,  after  several  days  a  small  part  is  found  to  have  com- 
bined to  form  water.  At  518°,  hours  are  required  before  the  union 
is  complete.  At  700°  the  combination  is  almost  instantaneous. 
Hence  contact  with  a  body  at  a  bright-red  heat  is  required  actually 
to  explode  the  mixture. 

Finely  divided  platinum,  when  held  in  the  cold  mixture,  hastens 
the  union  (otherwise  vanishingly  slow)  in  the  part  of  the  gases  hi 
contact  with  it.  The  heat  of  the  union  raises,  the  temperature  of 
the  platinum  and  of  neighboring  portions  of  the  gas  and  causes 
explosion  of  the  mass.  The  platinum  is  simply  a  contact  agent 
(p.  29)  and  remains  itself  unaffected. 

Hydrogen  unites  directly  with  a  minority  only  of  the  simple 
substances.  It  combines  rapidly  with  oxygen,  chlorine,  fluorine, 
and  lithium,  and  more  slowly  with  a  few  others. 

Hydrogen  acts  also  upon  some  of  the  compounds  of  metals  with 
oxygen  or  chlorine.  Thus,  when  any  one  of  the  oxides  of  iron  is 
heated  in  a  tube  through  which  hydrogen  flows,  the  latter  com- 
bines with  the  oxygen  to  form  water,  and  the  metal  is  liberated. 
The  skeleton  equation  (p.  51)  is:  Fe304  +  H  -»  H2O  +  Fe.  We 
then  reason  that  Fe3  will  give  3Fe.  Since  all  the  oxygen  is  removed 
from  the  compound,  O4  will  give  4H2O.  To  produce  this,  8H  is 
required.  Hence: 

Fe304  +  8H  ->  4H20  +  3Fe. 

This  interaction  is  classed  as  a  displacement.  In  describing  it  the 
chemist  would  also  say  that  the  hydrogen  has  been  oxidized  and 
that  the  oxide  of  the  metal  has  been  reduced  (pp.  36-37). 

The  Order  of  Activity  of  the  Metals.  —  We  employ  metals 
so  frequently  in  chemistry,  that  we  must  at  once  become  familiar 
with  the  key  to  the  main  differences  in  their  behavior.  The 
order  of  their  activity  explains  these  differences,  as  well  as  many 
other  facts.  In  the  adjoining  list,  the  most  active  metals  are 
at  the  top.  Hydrogen  is  not  a  metal,  but  is  included  because 
chemically  it  resembles  the  metals.  All  the  metals  above  hydrogen 
displace  this  element  from  dilute  acids  (and  from  water),  while 
those  below  it  do  not. 


60 


COLLEGE   CHEMISTRY 


The  first  displaces  the  hydrogen  from  water  violently,  the  second 
less  vigorously.  Magnesium  barely  acts  on  boiling  water,  but,  like 
iron,  acts  on  superheated  steam.  Zinc  liberates  hy- 
drogen with  reasonable  vigor  from  dilute  acids,  lead 
rather  feebly,  and  copper  and  those  following  not  at 
all. 

Other  facts  are  explained  by  the  table.  Thus,  when 
the  metals  are  heated  in  pure  oxygen,  the  last  two 
do  not  combine.  Those  above  silver  do  unite  with 
oxygen  —  mercury  rather  slowly  and  the  others  more 
and  more  energetically  as  we  ascend  the  list.  Again, 
if  we  take  the  oxides  of  the  metals,  we  find  that  those 
of  the  metals  up  to  and  including  mercury  lose  all 
their  oxygen  when. heated.  If  we  heat  the  oxides,  and 
lead  hydrogen  over  them,  the  oxygen  is  easily  removed 
from  all  the  oxides  up  to  and  including  those  of  iron, 
leaving  in  each  case  the  metal.  Thus,  in  general, 
the  more  active  metals  form  the  most  stable  com- 
pounds. 

The  metals  following  hydrogen  are  the  ones  which 
are  found  in  nature  in  large  amounts  in  the  free  con- 
dition. 


ORDER  OF 

ACTIVITY. 

METALS 

Potassium 

Sodium, 

Calcium 

Magnesium 

Aluminium 

Manganese 

Zinc 

Chromium 

Iron 

Nickel 

Tin 

Lead 

Hydrogen 

Copper 

Bismuth 

Antimony 

Mercury 

Silver 

Platinum 

Gold 


Exercises.  —  1.  Make  equations  for  reactions  in  which  hydro- 
gen is  liberated  by  the  action  of:  (a)  hydrochloric  acid  and  mag- 
nesium giving  MgCl2,  (6)  steam  and  zinc  giving  ZnO. 

2.  Make  an  equation  for  the  action  of  heat  on  manganese 
dioxide  Mn02  giving  oxygen  and  Mn304. 


CHAPTER  VI 
VALENCE.     CALCULATIONS 

Equivalence  and  Valence.  —  If  the  equations  showing  dis- 
placement of  hydrogen  by  a  metal  be  now  re-examined,  a  peculiar- 
ity will  be  observed  which  we  have  thus  far  omitted  to^note.  When 
sodium  (p.  50)  and  calcium  (p.  51)  act  upon  water,  one  atomic 
weight  (or  atom)  of  the  former  displaces  one  atomic  weight  of 
hydrogen,  but  one  atomic  weight  of  the  latter  displaces  twice  as 
much  hydrogen.  Again,  one  atom  of  zinc  (p.  53)  displaces  two 
atoms  of  hydrogen,  but  one  atom  of  aluminium  displaces  three. 
Assuming,  for  simplicity,  that  we  allow  three  of  these  metals  all 
to  act  upon  dilute  hydrochloric  acid,  the  equations  are: 

Na  +  HC1    ->  NaCl  +  H, 
Ca  +  2HC1  -»  CaCl2  +  2H. 
Al  +  3HC1  ->  A1C13  +  '3H. 

Interpreting  this,  we  perceive  that  the  atom  of  aluminium,  for 
example,  displaces  3H,  because  it  is  able  to  combine  with  3C7,  and 
so  incidentally  liberates  the  hydrogen  formerly  united  with  3C1. 
The  atom  of  sodium,  however,  can  unite  with  only  1C1,  and  so 
releases  only  1H.  Now  this  is  not  a  rule  confined  to  these  re- 
actions, but  represents  a  general  chemical  property  of  the  atomic 
weight  of  each  element,  and  a  property  which  we  shall  find  most 
useful. 

The  atom  of  aluminium  releases  3H  because  it  can  take  the 
place  of  three  atoms  of  hydrogen  in  chemical  combination  (and 
hold  3C1).  The  atomic  weight  of  aluminium  is  said  to  be  equiva- 
lent to  (equal  in  chemical  value  to)  three  atomic  weights  of  hydro- 
gen. Since  it  combines  with  3  atomic  weights  of  chlorine,  it  is  also 
considered  to  be  equi-valent  to  3  atomic  weights  of  this  element. 

The  chemical  property  referred  to  is  called  valence.  The  valence 
of  an  atomic  weight  of  hydrogen  or  of  chlorine  is  the  unit.  An 
atomic  weight  of  sodium  is  said  to  be  univalent,  one  of  calcium 

61 


62  COLLEGE   CHEMISTRY 

bivalent,  one  of  aluminium  trivalent.  The  formula  H20  shows  the 
atomic  weight  of  oxygen  to  be  bivalent,  because  it  unites  with  two 
atomic  weights  of  hydrogen.  Apparently,  the  atomic  weight  (or 
atom)  of  each  element  has  a  fixed  capacity  for  combining  with  not 
more  than  a  certain  number  of  atomic  weights  (or  atoms)  of  other 
elements. 

Marking  the  Valence.  —  Until  we  have  become  familiar  with 
the  valence  of  each  element,  it  is  advisable  to  mark  the  valences 
in  a  special  way:  Na1,  Can,  Alm,  Ou,  Zn11,  Cl1. 

As  we  should  expect,  a  bivalent  atom  can  combine  with  two 
univalent  atoms,  or  with  one  bivalent  atom,  and  so  forth.  Thus 
we  have  the  'compounds  of  oxygen:  Na^O11,  Ca1^)11,  A12III0311, 
ZnnOn,  CVO11. 

The  rule  is  that  the  quantities  of  two  elements  which  combine 
must  have  equal  total  combining  capacities  —  i.e.,  identical  total 
valence.  Thus,  Ca11  has  the  valence  two,  and  so  does  On.  Again, 
Al2m  has  a  total  valence  of  2  X  3  (  =  6)  and  so  has  O3n  (3x2  =  6). 

Frequently  the  valence  is  marked  by  means  of  lines,  the  num- 
ber of  lines  pointing  towards  a  symbol  indicating  the  valence  of 
the  atom  it  represents: 

/Cl  /Cl 

Na-Cl       Ca  Ca  =  O      A1-C1      O  =  A1-O-A1  =  O 

XC1  XC1 

Definition.  —  The  valence  of  an  element  is  a  number  repre- 
senting the  capacity  of  its  atomic  weight  to  combine  with,  or  dis- 
place, atomic  weights  of  other  elements,  the  unit  of  such  capacity 
being  that  of  one  atomic  weight  of  hydrogen  or  chlorine. 

Valence  of  Radicals.  —  What  we  have  said  applies  to  com- 
pounds of  not  more  than  two  elements  —  so  called  binary 
compounds.  We  cannot  with  certainty  tell  the  valences  in  a 
compound  of  three  or  more  elements,  like  H2SO4.  But  we  have 
seen  that  the  acids  behave  as  if  composed  of  two  radicals:  H(C1), 
H2(S04),  that  is,  of  two  groups  which  move  as  wholes  in  chemical 
reactions.  Hence  we  can  assign  a  valence  to  a  compound  radical 
as  a  whole.  Thus,  (S04)ir  is  evidently  bivalent,  as  a  whole,  be- 
cause it  is  united  with  2H1.  Na(OH)  and  Ca(OH)2  show  the 
radical  hydroxyl  (OH)  to  be  univalent. 


VALENCE.       CALCULATIONS  63 

It  is  to  preserve  the  identity  of  the  radicals,  and  to  make  them 
easily  recognizable,  that  we  write  them  in  brackets  and  place  the 
coefficient  outside,  as  Ca(OH)2  and  A12(S04)3,  instead  of  using  the 
forms  CaO2H2,  Al2S3Oi2,  and  so  forth.  In  fact,  substances  which 
commonly  interact  as  if  the  radicals  were  single  elements,  we  re- 
gard as  binary  compounds. 

In  writing  formulae  of  inorganic  compounds  we  usually  place 
the  positive  radical  (p.  53)  in  front  and  the  negative  radical  after  it. 

Use  in  Making  Formulse  and  Equations.  —  The  chief  use 
of  the  conception  of  valence  is  the  very  practical  one  of  enabling 
us  to  write  formulae.  In  making  equations  we  constantly  need  to 
know  whether  the  chloride  of  an  element,  say  magnesium,  is 
MgCl,  or  MgCl2,  or  MgCl3,  or  MgCU,  etc.,  and  whether  its  sul- 
phate is  MgS04,  or  Mg2SO4,  or  some  other  combination  of  the 
symbols.  To  answer  questions  like  this  it  is  not  necessary  to 
know  the  formula  of  every  compound  of  each  element;  the  ap- 
parent disorder  of  these  numbers  can  be  reduced  to  rule,  and  the 
reader  should  endeavor  thoroughly  to  master  the  rule  before  going 
farther. 

Thus,  suppose  that  we  require  the  formula  of  aluminium  hy- 
droxide. Up  to  this  point,  we  should  have  been  compelled  to 
look  for  it  in  a  book.  And  if,  later,  we  needed  the  formula  of 
aluminium  sulphate,  we  should  have  had  to  look  that  up,  sepa- 
rately, also.  But  now,  all  we  need  is  to  know  the  valence  of 
aluminium  Alm,  of  the  hydroxyl  radical  (OH)1  and  of  the  sul- 
phate radical  (804) n.  Making  the  total  valences  in  the  two 
halves  of  each  compound  alike,  we  write  the  formulae  A1III(OH)3I, 
A12III(S04)311. 

The  reader  must  make  a  special  effort  to  note  the  valences  of 
each  element  and  radical,  and  always  to  use  them  in  making 
formulae.  If  a  formula  is  written  from  memory,  the  valences 
must  be  checked,  to  make  sure  that  the  formula  is  correct. 

How  to  Learn  the  Valence  of  an  Element.  —  To  find  out 
the  valence  of  an  element,  we  must  obtain  the  formula  of  one  simple 
compound  of  the  element,  containing  another  element  of  known 
valence.  Thus,  what  is  the  valence  of  carbon?  Its  oxide  is  C02. 
The  total  valence  of  oxygen  here  is  2  X  2  =  4.  Carbon  CIV  is 


64  COLLEGE   CHEMISTRY 

therefore  quadrivalent.  Hence  its  chloride  must  be  C^CV  (carbon 
tetrachloride),  and  it  should  give  a  compound  with  hydrogen  C17!^1 
(methane,  composing  a  large  part  of  natural  gas).  When  carbon 
combines  with  a  trivalent  element,  equi-valent  amounts  of  each 
element  must  be  used,  as  in  Al^Cs^  (aluminium  carbide),  where 
AU111  and  C^  contain  3  X  4,  or  12  units  of  valence  each. 

The  chemist  does  not  memorize  the  valences  themselves;  he 
recovers  the  valence  of  an  element  or  radical,  when  needed,  by  recall- 
ing the  formula  of  a  substance  containing  this  element  or  radical 
in  combination  with  a  more  familiar  element  or  radical,  such  as  CP 
orH1. 

Elements  with  More  than  One  Valence.  —  The  rule  of 
valence  is  somewhat  complicated  by  the  fact  that  many  elements 
show  more  than  one  valence.  In  other  words,  the  combining 
capacity  of  an  atomic  weight  of  such  an  element  may  have  two 
(or  even  more)  values,  according  to  the  conditions  under  which 
the  action  takes  place. 

Thus,  we  have  encountered  two  chlorides  of  iron,  ferrous  chloride 
FenCl2I  and  ferric  chloride  Fe™^1.  We  have,  in  fact,  two  com- 
plete series  of  compounds  of  iron,  such  as: 

Bivalent  (Ferrous):      FeCl2,    FeO,      FeSO4. 
Trivalent  (Ferric):       FeCl3,   Fe203,   Fe2(S04)3. 

When  an  element  forms  two  such  series  of  compounds,  we  always 
call  particular  attention  to  the  fact. 

Exceptional  Valences.  —  Some  elements  show  an  exceptional 
valence  in  one  compound.  The  valences  shown  in  series  of  com- 
pounds are  the  important  ones,  and  the  exceptions  need  not 
particularly  concern  us.  Thus,  in  addition  to  the  oxides  FeO  and 
Fe^Os,  iron  gives  the  magnetic  oxide  Fe3O4,  where  the  valence  of 
iron  appears  not  to  be  a  whole  number,  but  f  or  2f .  Hence  the 
valence  is  made  regular  by  supposing  the  oxide  to  be  a  compound 
of  the  other  two  oxides,  as  if  the  formula  were  FenO,Fe2lnO3. 

Nomenclature.  —  The  names  of  compounds  containing  only 
two  elements  (the  true  binary  compounds)  end  in  ide.  Such  are 
the  oxides,  as  ferric  oxide  FeaOs;  the  carbides,  as  aluminium  car- 


VALENCE.      CALCULATIONS  65 

bide  AkC3;  the  chlorides,  as  sodium  chloride  NaCl;  the  sulphides, 
as  ferrous  sulphide  FeS,  etc. 

When  an  element  forms  two  (or  more)  compounds  with  another 
element,  they  are  frequently  distinguished  thus:  carbon  dioxide 
C02,  carbon  monoxide  CO;  phosphorus  peroxide  P205,  phosphorus 
dioxide  P20s. 

To  distinguish  two  compounds  of  the  same  elements,  another 
plan  is  also  used:  ferrous  chloride  FeCl2,  ferric  chloride  FeCls; 
mercurows  oxide  Hg20,  mercuric  oxide  HgO.  The  suffix  OILS  indi- 
cates that  the  metal  is  combined  with  the  smaller  proportion  of 
the  negative  element,  and  ic  that  it  is  combined  with  the  larger 
proportion. 

The  tendency  —  although  not  a  universal  rule  —  is  to  use  the 
latter  plan  with  compounds  containing  a  metal  and  the  former 
with  compounds  containing  only  non-metals. 

Equivalent  Weights.  —  In  the  foregoing  discussion  of  valence, 
we  have  more  than  once  used  the  word  "  equivalent."  For 
example  (p.  61),  it  was  stated  that  the  atomic  weight  of  aluminium 
is  equivalent  to  three  atomic  weights  of  hydrogen,  because  it  dis- 
places them,  and  to  three  atomic  weights  of  chlorine,  because  it 
combines  with  that  number: 

Al     +        3HC1      ->  AlClj  +        3H 

Weights:     27.1         3  X  36.468       27.1  +  3  X  35.46      3  X  1.008 

Now  chemists  often  view  this  from  the  other  direction,  and  say 
that  1.008  g.  of  hydrogen  are  displaced  by  9.03  g.  of  aluminium 
(one-third  of  the  atomic  weight)  and  that  35.46  g.  of  chlorine 
combine  with  only  9.03  g.  of  aluminium.  When  taking  this  view, 
they  refer  to  the  weight  of  an  element  displacing  one  atomic  weight 
of  hydrogen,  or  combining  with  one  atomic  weight  of  chlorine  (or  of 
any  other  univalent  element)  as  the  equivalent  weight  of  that  ele- 
ment. The  equivalent  weight  of  aluminium  is  therefore  9.03  and 
that  of  calcium  Ca11  20  (one-half  the  atomic  weight)  and  that  of 
sodium  Na1  23  (the  atomic  weight). 

It  will  be  seen  that  the  equivalent  weight  can  always  be  found 
by  a  quantitative  experiment.  It  is  also  evident  that  it  is  equal 
to  the  atomic  weight  divided  by  the  valence.  It  is  likewise  clear 
that  the  equivalent  weight  of  an  element,  multiplied  by  the  valence 


66  COLLEGE   CHEMISTRY 

of  that  element,  is  equal  to  the  atomic  weight.  The  conception 
of  equivalent  weights  finds  application  in  several  connections  in 
chemistry  (see  Normal  Solutions  and  Faraday's  Law). 

CALCULATIONS 

As  we  have  seen  (p.  44),  the  formula  represents  the  composition 
of  a  substance,  using  the  atomic  weights  as  the  units.  We  have 
learned  how  the  formula  is  calculated  from  measurements  made 
in  an  experiment  (p.  45).  We  may  now  take  up  some  of  the  ways 
of  using  the  information  contained  in  a  formula. 

Composition  from    the    Formula.     Formula-Weight. — 

To  learn  the  composition  of  a  substance,  such  as  potassium 
chlorate,  KC103,  from  its  formula,  we  look  up  the  values  of  the 
atomic  weights  (inside  rear  cover).  We  find  K  =  39.1  parts  of 
potassium,  Cl  =  35.46  parts  of  chlorine,  and  03  =  3  X  16.00  or 
48  parts  of  oxygen.  The  proportions,  in  order,  are  therefore: 
39.1  :  35.46  :  48. 

What  is  the -proportion  of  oxygen  to  potassium  and  chlorine, 
together?  It  is  48  :  39.1  +  35.46,  or  48  :  74.56,  or  1  :  1.55. 

We  require  a  name  for  the  sum  of  the  weights  of  the  constituents 
indicated  in  the  formula.  This  is  called  the  formula-weight. 
Thus,  for  potassium  chlorate,  it  is  39.1  +  35.46  +  48,  or  122.56. 

To  Find  the  Percentage  Composition.  —  In  potassium 
chlorate  the  proportions  are  39.1  of  potassium,  35.46  of  chlorine, 
and  48  of  oxygen  or  a  total  of  122.56.  In  one  hundred  parts,  the 

potassium  is  ^^  X  100,  or  31.9%;  the  chlorine  ^^  X  100, 

1ZZ.OO  l^JZ.OO 

or  28.9%;  and  the  oxygen  ^^  X  100,  or  39.1%. 

Stated  in  terms  of  the  rule  of  proportion,  we  have,  for  the  potas- 
sium, 122.56  :  39.1  ::  100  :  x,  where  x  is  the  percentage  of  potassium. 

Calculations  by  Use  of  Equations.  —  We  frequently  wish 
to  know  what  weight  of  a  product  can  be  obtained  from  a  given 
weight  of  the  necessary  materials.  For  example,  what  weight  of 
ferrous  sulphide  can  be  made  with  100  g.  of  iron?  It  is  under- 
stood that  the  necessary  sulphur  is  available. 


CALCULATIONS  67 

To  avoid  the  blunders  which  are  easily  made,  observe  strictly 
the  following  rules: 

1.  Write  down  the  equation: 

Fe  +  S  -»  FeS. 

2.  Place  under  each  formula  the  weight  it  represents: 

Fe     +      S      ->    FeS 
55.84        32.06        87.90 

3.  Read  this  expanded  equation.     In  this  case  it  reads:    55.84 
parts  of  iron  combine  with  32.06  parts  of  sulphur  to  give  87.90 
parts  of  ferrous  sulphide. 

4.  Re-read  the   original   problem:     "What   weight   of  ferrous 
sulphide  can  be  made  with  100  g.  of  iron? "     Having  done  this, 
place  the  amount  given  in  the  problem  (100  g.  of  iron)  under  the 
formula  of  the  substance  in  question.     Then  notice  what  the  prob- 
lem asks  ("what  weight  of  ferrous  sulphide")  and  place  an  x  under 
the  formula  of  that  substance: 

Fe     +       S     ->    FeS 
55.84        32.06        87.90 
100  g.  x 

5.  Read  the  problem  as  now  tabulated:  55.84  g.  of  iron  give  87.90 
g.  of  ferrous  sulphide,  therefore  100  g.  of  iron  will  give  x  g.  of 
ferrous  sulphide. 

6.  State  the  proportion  in  this  order  (or,  see  below). 

55.84  :  87.90  ::  100  :  x  (  =  157.4  g). 

If  the  tabulation  in  rule  4  has  been  prepared  correctly,  this  final 
statement  as  a  proportion  is  purely  mechanical.  It  will  be  noted 
that  only  two  of  the  three  quantities  given  in  the  expanded  equa- 
tion were  actually  used. 

6a.   Alternative  Method.     At  the  sixth  step,  we  may  also  say: 
If  55.84  g.  of  iron  give  87.90  g.  of  ferrous  sulphide,  1  g.  of  iron  will 

give  |^|5  g.  (=  1.574  g.)  of  ferrous  sulphide.     Then,  if  1  g.  of 

iron  gives  1.574  g.  of  ferrous  sulphide,  the  100  g.  of  iron  will  give 
100  X  1.574  g.  (  =  157.4  g.)  of  ferrous  sulphide. 


68  COLLEGE    CHEMISTRY 

Warnings.  —  In  solving  the  exercises  at  the  end  of  the  chap- 
ter, beware  of  three  kinds  of  mistakes,  which  are  commonly  made. 

1.  Do  not  read  the  problem  carelessly  and  make  the  equation 
backwards,  that  is,  with  the  sides  reversed.     Focus  attention  first 
on  the  exact  chemical  change  involved. 

2.  Do  not  speak,  or  think  of  the  symbols  Fe  and  S  as  standing 
for  "1  part"  of  iron  or  sulphur.     They  stand  for  1  chemical  unit, 
or  atomic  weight,  or  atom,  in  each  case,  that  is,  for  "55.84  parts  " 
and  "32.06  parts,"  respectively. 

3.  Follow  the  rules  laid  down  above.     The  chemist  follows 
these  rules.     The  beginner  always  thinks  he  can  do  without  them, 
and  he  fails  in  consequence.     Writing  the  equation  in  expanded 
form  (rule  4)  and  reading  the  problem  into  it  (rule  5)  are  abso- 
lutely essential  steps. 

Another  Example.  —  What  weight  of  hydrogen  is  required 
to  reduce  45  g.  of  magnetic  oxide  of  iron  to  metallic  iron? 
Following  the  rules,  as  before,.  we  reach  the  expanded  equation: 

Fe304  +      4H2      ->      3Fe  +          4H20. 

3  X  55.84  +  4  X  16          8  X  1.008         3  X  55.84         4(2  X  1.008  +  16) 

167.52  +  64  8.064  167.52  4X18.016 

231.52  8.064  167.52  72.064 

45  g.  x 

Observe  that  the  atomic  weights  are  multiplied  by  the  sub- 
numbers,  so  that,  for  example,  Fe3  =  3  X  55.84.  Observe  also 
that  the  formula-weights  are  multiplied  by  the  coefficients,  when 
such  occur  in  front  of  the  formulae,  so  that,  for  example,  4H20  =  4 
X  18.016. 

The  proportion  231.52  :  8.064  ::  45  :  x  (  =  1.57)  supplies  the 
answer,  1.57  grams  of  hydrogen. 

Using  the  alternative  plan:   If  231.52  g.  of  magnetic  oxide  "are 

reduced  by  8.064  g.  of  hydrogen,  1  g.  will  be  reduced  by  8'^64  g. 


(=  0.035  g.)  of  hydrogen.  Hence,  if  1  g.  of  magnetic  oxide  is 
reduced  by  0.035  g.  of  hydrogen,  45  g.  will  be  reduced  by 
45  X  0.035  g.  (=1.57  g.)  of  hydrogen. 

Exercises.  —  1.   What  are  the  valences  of  the  negative  radicals 
of  phosphoric  acid  H3PO4,  and  of  acetic  acid  (p.  52)?    What  must 


CALCULATIONS  69 

be  the  formulae  of  calcium  phosphate,  cupric  acetate  (Cu11),  alumin- 
ium phosphate,  ferrous  carbonate  (C03n),  ferrous  sulphate, 
cupric  chloride? 

2.  What  is  the  valence  of  phosphorus  in  phosphorus  pentoxide 
(p.  32)?     What  must  be  the  formulae  of,  (a)  the  corresponding 
chloride  and  sulphide  of  phosphorus,  and  (6)  of  aluminium  oxide? 

3.  What  are  the  valences  of  the  elements  in  the  following:  LiH, 
NH3,  SeH2,  BN? 

4.  What  are  the  valences  of  the  metals  and  radicals  in  the  fol- 
lowing:   HNO3,  Pb(NO3)2,   Ce(S04)2,   KC1,   KMn04   (potassium 
permanganate)?     Name  all  the  substances  in  3  and  4. 

5.  Make  equations  to  represent,  (a)  the  reduction  of  lead  dioxide 
(Pb02)  by  hydrogen,  (b)  the  actions  of  aluminium  upon  cold  water 
and,  (c)  upon  steam  at  a  red  heat. 

6.  What  weight  of  mercury  is  obtained  from  120  g.  of  mercuric 
oxide  HgO? 

7.  What  weight  of  mercuric  oxide  will  furnish  20  g.  of  oxygen? 

8.  What  weight  of  Fe203  may  be  obtained  from  10  g.  of  oxygen? 

9.  How  much  silver  is  contained  in  100  g.  of  an  impure  specimen 
of  silver  chloride  AgCl  which  is  33  per  cent  sand? 

10.  What  are  the  percentage  compositions  of  cerium  sulphate 
Ce(S04)2,  phosphorus  pentachloride  PCls,  and  ammonium  chloride 


11.   What  weight  of  hydrogen  is  required  to  reduce  100  g.  of 
ferric  chloride  FeCl3  to  ferrous  chloride  FeCl2? 


CHAPTER  VII 

MEASUREMENT  OF    QUANTITY    IN    GASES.     RELATIONS 
BETWEEN    STRUCTURE    AND    BEHAVIOR    OF    MATTER 

A  SPECIMEN  of  a  gas,  like  a  specimen  of  a  solid  or  a  liquid,  may  be 
weighed,  but  it  is  usually  easier  to  determine  the  quantity  of  the 
gas  by  (1)  measuring  its  volume,  and  at  the  same  time  (2)  noting 
its  temperature  on  a  thermometer  suspended  in  it  or  close  to  it,  and 
(3)  ascertaining  the  pressure  which  it  exercises. 

The  Measurement  of  the  Pressure  of  a  Gas.  —  In  almost  all 
cases  the  easiest  way  to  take  account  of  the  pressure  of  a  gas  is  to 
place  it  in  an  apparatus  so  constructed  that  one 
boundary  of  the  volume  is  a  liquid.  The  apparatus 
is  then  so  adjusted  that  the  surface  of  the  liquid  in 
contact  with  the  gas  in  the  closed  tube  (Fig.  33)  is  at 
the  same  level  as  the  free  surface  of  the  liquid  which  is 
exposed  to  the  atmosphere.  The  equality  in  the  levels 
of  the  liquids  is  then  a  guarantee  that  the  specimen 
of  gas  and  the  atmosphere  are  exercising  equal  pres- 
sures on  the  liquid.  At  this  stage  the  volume  of 
the  gas  is  measured,  by  reading  the  graduation  (not 
shown)  on  the  tube.  Simultaneously  the  pressure  of  the  atmos- 
phere and,  therefore,  of  the  gas,  is  ascertained  by  reading  the 
barometer. 

The  barometer  (Fig.  34)  consists  of  a  bent  tube  containing  mer- 
cury. The  short  limb  (to  the  left)  is  open  and  the  pressure  of  the 
atmosphere  is  exercised  on  the  surface  of  the  mercury  there.  The 
longer  limb  (to  the  right)  is  closed  at  the  top  and  in  it  there  is  no 
gas  above  the  mercury.  When  the  tube  is  inclined,  the  surface  of 
the  mercury  in  the  longer  limb  endeavors  to  retain  the  same  verti- 
cal height  above  the  lower  surface  and  consequently  rises  and,  with 
sufficient  inclination,  will  reach  entirely  to  the  top  of  the  tube. 
The  downward  pressure  of  the  mercury  on  the  right,  above  the 

70 


THE   MEASUREMENT   OF   QUANTITY   IN    GASES 


71 


dotted  line,  is  exactly  equal  to  the  pressure  of  the  atmosphere  on  the 
free  surface  of  the  mercury  at  the  same  level.  The  amount  of  the 
latter  pressure  is  proportional  to  the  length  of  the  column  of  mer- 
cury above  the  dotted  line.  Hence,  reading  the 
height  at  which  the  mercury  stands  above  the  free 
surface  gives  us  a  measure  of  the  pressure  of  the 
atmosphere  and  of  any  specimen  of  gas  which  is 
at  the  same  pressure. 

This  is  called  the  uncorrected  reading.  It  is 
immediately  reduced  to  the  reading  which  would 
have  been  made  if  the  barometer  and  its  ^mercury 
had  been  at  0°  (corrected  reading),  by  noting  the 
temperature  on  the  adjacent  thermometer  and 
subtracting  from  the  uncorrected  reading  the 
necessary  correction  (Table  of  Corrections,  C, 
Fig.  34). 

\  For  example:  the  volume  of  gas,  after  adjust- 
/  ment  to  atmospheric  pressure,  is  200  c.c.  and  its 
temperature  17°.  The  uncorrected  barometric 
reading  is  744  mm.  with  the  barometer  (perhaps 
in  a  different  room  from  the  gas)  at  15°.  The 
correction  is  —2.0  mm.  The  corrected  reading  is 
therefore  742  mm. 

Fia.  34. 

Correction    of    the    Volume    to    760     mm.    Pressure.  — 

Finally,  since  the  atmospheric  pressure  varies  from  day  to  day, 
the  volume  at  the  observed  pressure  is  corrected  to  that  which  the 
same  quantity  of  gas  would  have  occupied  at  the  standard  pressure 
of  760  mm.  of  mercury.  By  careful  measurements,  Boyle  (1660) 
found  that  the  volume  occupied  by  the  same  sample  of  any  gas 
varies  inversely  with  the  pressure. 

The  illustration  just  given  will  show  how  this  additional  correc- 
tion is  applied.  There  were  200  c.c.  of  the  gas  at  17°  and  742  mm. 
(corr.).  The  question  is:  What  would  be  the  volume  of  this 
amount  of  gas  at  760  mm.?  At  this  new  pressure  (760  mm.),  which 
is  greater  than  the  old  pressure  (742  mm.),  the  volume  will  become 
less.  Hence  we  change  the  volume  in  the  proportion  of  these 
pressures,  placing  the  smaller  number  in  the  numerator,  so  as  to  get  a 
smaller  volume  as  the  answer:  200  X  J|£  =  volume  at  760  mm., 


72  COLLEGE    CHEMISTRY 

=  195.3  c.c.  If  we  wished  to  convert  100  c.c.  at  775  mm.  to  760 
mm.,  we  should  reason  that  the  new  pressure  was  smaller,  and  the 
volume  would  become  greater,  and  should  therefore  place  the 
larger  number  (775)  in  the  numerator  so  as  to  get  a  larger  volume 
for  the  answer. 

The  Correction  of  the  Volume  of  a  Gas  for  Temperature. — 

The  same  sample  of  gas  will  occupy,  when  heated,  a  larger  volume, 
and  when  cooled,  a  smaller  volume  than  before.  The  change  in 
volume  for  each  degree  Centigrade  is  •$$?  of  the  volume  of  the  same 
sample  at  0°.  To  simplify  the  calculation  we  begin  by  converting 
the  temperature  to  the  absolute  scale  by  adding  273°  to  each 
temperature.  The  volumes  assumed  by  a  sample  of  gas  at  different 
temperatures,  the  pressure  remaining  constant,  are  in  the  same  pro- 
portion as  the  corresponding  absolute  temperatures  (Charles,  1787). 
If  the  volume  remains  constant,  then  the  pressure  changes  in  the 
same  proportion. 

In  the  illustration  used  above,  there  were  200  c.c.  of  gas  at  17°, 
and  it  is  required  to  know  the  volume  at  0°.  We  add  273  algebra- 
ically to  each  temperature,  and  the  question  becomes:  There  are 
200  c.c.  of  gas  at  290°  Abs.,  what  will  be  its  volume  at  273°  Abs.? 
The  volume  changes  in  the  direct  ratio  of  the  absolute  temperatures. 
The  new  temperature  is  lower  than  the  old,  and  the  new  volume  will 
therefore  be  smaller  than  the  old.  Then  200  X  JJ$  =  volume  at 
0°  (273°  Abs.)  =  188.3  c.c. 

The  above  laws  are  usually  applied  to  any  example  simulta- 
neously. Thus,  200  c.c.  of  gas  at  742  mm.  pressure  (corr.)  and  17° 
become  200  X  ft}  X  H8  =  183.8  c.c.  at  0°  and  760  mm. 

Mixed  Gases:  Aqueous  Tension.  —  Two  gases  at  the  same 
temperature,  provided  they  do  not  interact  chemically,  do  not  in- 
terfere with  each  other's  pressures  when  mixed.  Thus,  if  they  are 
forced  into  the  same  volume,  the  pressure  of  the  mixture  is  equal 
to  the  sum  of  those  of  the  components  (Dalton's  law,  1807).  The 
gases  are  therefore  still  thought  of  individually,  and  the  share 
which  each  gas  has  in  the  total  pressure  is  called  its  partial  pres- 
sure. This,  like  any  other  gaseous  pressure,  is  proportional  to 
the  concentration  of  the  particular  gas  in  the  mixture. 

For  example,  a  gas  measured  over  water  contains  water  vapor. 


THE   MEASUREMENT   OF   QUANTITY   IN   GASES  73 

The  partial  pressure  of  this,  called  aqueous  tension  (q.v.),  which 
is  definite  for  each  temperature,  must  be  subtracted  from  the  total 
pressure.  The  remainder  is  the  partial  pressure  of  the  gas  being 
measured,  and  this  remainder  is  used  as  the  pressure  of  this  gas  in 
any  calculation.  Thus,  in  a  gas  measured  over  water  at  22°,  the 
total  pressure  includes  19.7  mm.  pressure  of  water  vapor  (the 
aqueous  tension  at  22°,  see  Appendix  IV).  Hence  150  c.c.  of  gas 
over  water  at  22°  and  750  mm.  is  the  same  in  amount  as  150  c.c.  of 
the  same  gas  in  dry  condition  at  22°  and  730.3  mm.  (there  being 
simply  150  c.c.  of  water  vapor  at  19.7  mm.  mixed  with  it).  To 
obtain  the  volume  of  dry  gas  at  0°  and  760  mm.  we  have  the  ex- 
pression 150  X  fit  X 


Densities  of  Gases.*  —  The  density  of  a  gas  is  the  mass  of  1  c.c. 
of  the  gas  at  0°  and  760  mm.  pressure.  Sometimes  the  weight 
of  one  liter  (1000  c.c.)  is  called  the  density.  Often  the  relative 
weight  of  the  gas,  the  weight  of  an  equal  volume  of  air,  or  oxygen, 
or  hydrogen  being  taken  as  unity,  receives  the  same 
name. 

The  most  direct  method  of  measuring  the  density 
of  a  gas  is  to  employ  a  light  flask  of  125-150  c.c. 
capacity,  provided  with  a  rubber  stopper  and  stop- 
cock (Fig.  35).  By  means  of  an  air-pump  the  con- 
tents of  the  flask  are  removed,  and  it  is  weighed. 
This  gives  the  weight  of  the  empty  vessel.  The  gas, 
whose  density  is  to  be  ascertained,  is  then  admitted, 
and  care  is  taken  that  it  finally  fills  the  flask  at 
the  pressure  of  the  atmosphere.  The  flask  is  *"*- 35- 
closed  and  weighed  again.  The  increase  represents  the  weight  of 
the  gas.  At  the  same  time  the  temperature  and  barometric 
pressure  are  read.  The  volume  is  determined  by  displacing  the 
gas  once  more  from  the  flask,  filling  with  water,  and  weighing 
again.  The  difference  in  weight  between  the  empty  flask  and  the 
flask  full  of  water,  in  grams,  represents  the  volume  of  the  content 
of  the  flask  in  cubic  centimeters.  This  volume  is  reduced  to  0° 
and  760  mm.  by  the  rules  discussed  above,  and  we  have  then  a 
volume  of  the  gas  and  the  corresponding  weight. 

*  The  subjects  of  this  section  are  not  actually  used  until  Chapter  IX  (on 
Molar  Weights)  is  reached. 


74  COLLEGE    CHEMISTRY 

To  illustrate,  let  us  suppose  that  the  volume  of  the  flask  is  200  c.c. 
and  that  it  is  filled  with  oxygen  at  17°  and  742  mm.  The  weight 
of  the  gas  is  found  to  be  0.26  g.  We  ascertained  (p.  72)  by  calcula- 
tion that  at  0°  and  760  mm.  this  volume  would  be  183.8  c.c.  The 
weight  of  a  liter  is  given  by  the  proportion  183.8  :  0.26  ::  1000  :  x. 
Here  x  =  1.415  g.  When  the  operation  is  performed  carefully, 
and  the  weighing  carried  to  the  nearest  milligram  instead  of  the 
nearest  centigram,  a  result  more  nearly  approaching  the  accepted 
one  (1.429)  may  easily  be  reached. 

To  get  the  density  of  oxygen  referred  to  hydrogen  £s  unity,  we 
must  divide  the  answer  by  the  weight  of  a  liter  of  hydrogen 
(0.08987  g.).  In  the  above  case  the  quotient  is  15.74.  The  ac- 
cepted value  is  15.90.  The  density  referred  to  air  as  unity  is 
similarly  obtained  by  dividing  by  1.293,  the  weight  of  a  liter  of  air 
at  0°  and  760  mm.  pressure. 

By  using  a  modification  of  the  flask  just  described,  it  is  possible 
to  ascertain  the  weights  of  known  volumes  of  the  vapors  of  liquids 
and  solids.  A  temperature  sufficiently  high  to  vaporize  the  sub- 
stance must  be  employed.  The  volume  is  reduced  by  rule  to  0° 
and  760  mm.  and  the  density  (in  this  case  known  as  the  vapor 
density)  is  calculated  as  before.  The  reduction  to  0°  and  760  mm. 
pressure  by  rule  gives,  of  course,  a  fictitious  result.  The  vapor 
would  condense  to  the  liquid  form  before  0°  was  reached,  if  the 
cooling  were  actually  carried  out.  But  the  value  for  the  density 
as  it  would  be  at  0°  and  760  mm.  has  to  be  calculated  to  facilitate 
comparison  with  the  corresponding  values  for  other  substances. 
The  results  have  no  physical  significance,  but  are  highly  important 
to  the  chemist. 

RELATIONS  BETWEEN  THE  STRUCTURE  AND  BEHAVIOR  OF 
MATTER 

We  have  seen  that  matter  is  composed  of  minute  particles 
called  molecules.  Just  as  we  can  thoroughly  understand  the  be- 
havior of  a  watch  or  an  automobile  engine  only  if  we  know  the 
details  of  its  structure,  and  how  the  parts  work,  so  we  can  under- 
stand the  physical  and  chemical  behavior  of  matter  in  masses 
only  if  we  are  familiar  with  its  ultimate  mechanism.  Hence,  we 
must  now  take  up  the  structure  of  matter  in  its  three  states,  the 


STRUCTURE   AND   BEHAVIOR   OF   MATTER  75 

gaseous,  the  liquid,  and  the  solid.  In  doing  this,  we  shall  keep 
constantly  in  view  the  connection  between  the  molecular  relations 
and  the  general  behavior  of  the  matter. 

The  Molecular  Structure  of  Gases.  —  The  most  noticeable 
fact  about  gases  is  that  they  can  be  compressed  to  an  enormous 
extent.  Oxygen  at  760  mm.,  for  example,  can  be  reduced  by 
pressure  to  one  two-hundredth  of  its  volume,  or  even  less.  The 
compression  does  not  affect  the  individual  molecules,  and  there- 
fore does  not  diminish  the  volume  actually  occupied  by  the  oxygen, 
but  it  crowds  the  molecules  closer  together  and  diminishes  to  one 
two-hundredth  the  space  between  them.  Compressing  a  gas  is, 
in  fact,  mainly  compressing  the  empty  space  of  which  it  chiefly 
consists.  To  understand  what  follows,  the  reader  must  keep 
constantly  before  him  a  mental  image  of  a  jar  of  gas  as  consisting 
of  small  particles  separated  by  relatively  wide,  empty  spaces. 
The  molecules  are  in  rapid  motion  and  move  in  straight  lines,  ex- 
cepting when  they  strike  one  another  or  the  walls  of  the  vessel. 

The  Properties  of  Gases.  —  Let  us  now  note  the  more  obvious 
qualities  of  gases,  printing  in  italics  the  fact  concerning  a  mass  of 
gas  and  in  black  type  the  property  of  the  molecules  which  accounts 
for  the  fact. 

[  The  most  remarkable  thing  about  a  gas,  considering  the  loose- 
ness with  which  its  material  is  packed,  is  the  total  absence  in  it  of 
any  tendency  to  settling  or  subsidence.  Since  the  molecules  cannot 
be  at  rest  upon  one  another,  as  the  great  compressibility  shows,  we 
are  driven  to  conclude  that  they  are  widely  separated  from  one 
another,  and  that  they  occupy  the  space,  otherwise  a  complete 
vacuum,  by  constantly  moving  about  in  all  directions.  But  a 
moving  aggregate  of  particles  which  does  not  even  finally  settle 
must  be  in  perpetual  motion.  We  must,  therefore,  believe  the 
molecules  to  be  wholly  unlike  particles  of  matter  in  having  perfect 
elasticity,  in  consequence  of  which  they  undergo  no  loss  of  energy 
after  a  collision.  They  must  continually  strike  the  walls  of  the 
vessel  and  one  another  and  rebound,  yet  without  loss  of  motion. 
The  fact  that  each  gas  is  homogeneous,  efforts  to  sift  out  lighter  or 
heavier  samples  having  failed,  requires  the  supposition  that  all  the 
molecules  of  a  pure  gas  are  closely  alike. 


76  COLLEGE  'CHEMISTRY 

The  diffusibility  of  gases  is  due  to-  the  motion  of  the  molecules, 
and  their  permeability  to  the  space  available  to  receive  molecules 
of  another  gas.  These  two  modes  of  behavior  involve  no  additional 
molecular  properties.  The  word  "diffusion"  is  often  thought  to 
mean  the  property  of  a  given  mass  of  gas  in  virtue  of  which  another 
gas  can  mix  with  the  given  mass.  This  property  is  not  diffusibility 
but  permeability.  It  is  the  other  gas,  which  makes  its  way  into 
the  given  gas,  which  is  diffusing.  Diffusion  is  spontaneous 
motion  of  the  parts  of  a  gas  away  from  their  original  location. 
Unless  this  motion  is  into  an  empty  space,  the  diffusing  molecules 
must,  of  course,  move  into  another  body  of  gas.  In  the  case  of 
the  jars  of  hydrogen  and  air  (p.  57),  each  gas  moved  in  part  out  of 
its  original  jar  (diffused),  and  each  received  parts  of  the  other  gas 
into  its  jar  (was  permeated). 

Boyle9 s  Law  and  Charles9  Law.  —  Passing  now  to  Boyle's 
law  (p.  71),  the  thing  to  be  accounted  for  is  that  when  a  sample  of 
a  gas  diminishes  in  volume,  its  pressure  increases  in  the  same  pro- 
portion. Let  the  diagram  (Fig.  36)  represent  a  cylinder  with  a 
movable  piston,  upon  which  weights  may  be  placed  to  resist  the 
pressure.  Now  the  pressure  exercised  by  the  gas  under 
the  piston  cannot  be  like  the  pressure  of  the  hand  upon 
a  table,  since  we  have  just  assumed  that  the  particles 
are  not  even  approximately  at  rest,  and  the  spaces 
between  them  are  enormous  compared  with  the  size 
of  the  molecules  themselves.  The  gaseous  pressure 
must  therefore  be  attributed  to  the  colossal  hailstorm 
which  their  innumerable  impacts  upon  the  piston 
produce.  If  this  is  the  case,  the  compressing  of  a  gas 
must  consist  simply  in  moving  the  partition  down- 
wards, so  that  the  particles  as  they  fly  about  are  gradu- 
ally restricted  to  a  smaller  and  smaller  space.  Their  paths  become 
on  an  average  shorter  and  shorter.  Their  impacts  upon  the  walls 
become  more  and  more  frequent.  So  the  pressure  which  this 
causes  becomes  greater  and  greater,  and  is  proportional  to  the 
degree  of  crowding  (the  concentration)  of  the  molecules. 

There  are  two  other  points  to  be  added.  When  we  diminish  the 
volume  to  one-half,  we  find  from  experience  that  the  pressure  be- 
comes exactly,  or  almost  exactly,  twice  as  great.  This  must  mean 


STRUCTURE   AND    BEHAVIOR   OF  MATTER  77 

that,  although  the  particles  are  becoming  crowded,  they  do  not 
interfere  with  one  another's  motion,  excepting  of  course  where 
actual  collision  causes  a  rebound.  Only  in  the  absence  of  inter- 
ference would  doubling  the  number  of  molecules  per  unit  of  volume 
give  exactly  double  the  number  of  impacts  on  the  walls.  Hence 
the  molecules  must  have  practically  no  tendency  to  cohesion. 
Finally,  the  molecules  must  be  supposed  to  move  in  straight  lines 
between  collisions. 

A  baseless  idea,  that  the  molecules  of  a  gas  repel  one  another, 
still  lingers  in  some  quarters.  There  is  no  evidence  of  this.  The 
molecules  pay  almost  no  attention  to  one  another,  either  by 
attraction  or  repulsion. 

Boyle's  law  therefore  adds  four  more  details  concerning  molec- 
ular behavior,  namely,  that  the  impacts  of  the  particles  produce 
the  pressure,  that  the  crowding  of  the  molecules  represents  the  con- 
centration of  the  material  and  that  the  particles  move  in  straight  lines 
and  show  almost  no  cohesion,  since  pressure  and  concentration  are 
very  closely  proportional  to  one  another. 

How,  now,  can  we  account  for  Charles'  law  (p.  72),  according  to 
which  an  increase  in  pressure  (or  in  volume)  results  from  heating 
a  mass  of  rapidly  moving  molecules?  The  action  of  a  particle 
colliding  with  a  surface  is  measured  in  physics  in  terms  of  its 
mass  and  its  velocity.  It  is  evident  that  heating  a  cloud  of  mole- 
cules would  not  increase  the  mass  of  each,  and  it  must  therefore 
increase  the  velocity  of  each,  since  the  kinetic  energy  of  all  becomes 
greater. 

Avogadro's  Law.  —  The  identical  general  behavior  of  all 
kinds  of  gases  suggests  that  their  structures  may  be  all  alike. 
Avogadro  (1811),  the  professor  of  physics  in  Turin,  put  forward 
the  hypothesis  that  the  numbers  of  molecules  in  equal  volumes  of 
different  gases,  at  the  same  temperature  and  pressure,  might  be 
equal.  A  more  strict  study  of  the  properties  we  have  been  con- 
sidering, and  of  some  additional  facts,  has  since  shown  that  no 
other  conjecture  than  Avogadro's  would  be  consistent  with  them. 
Thus  it  is  now  accepted  as  a  fact,  and  is  known  as  Avogadro's  law. 
It  may  also  be  put  in  the  form:  At  the  same  temperature  and 
pressure,  the  molecular  concentration  of  all  kinds  of  gases  has  the 
same  value. 


78  COLLEGE    CHEMISTRY 

Diffusion.  —  The  law  of  diffusion  (p.  58)  harmonizes  with  the 
conceptions  of  molecular  structure  without  further  additions  to  the 
latter.  The  speed  of  the  hydrogen  molecule  at  room  temperature 
is  1840  meters  per  second.  The  masses  of  the  hydrogen  and  oxy- 
gen molecules  are  as  1  :  16,  and  the  speeds  of  diffusion  (p.  58)  as 
VI  :  Vl6,  or  1  :  4.  Hence  the  speed  of  the  oxygen  molecule  is 
one-fourth  of  1840,  or  460  m.  per  sec. 

Calculation  shows  the  activity  of  the  molecules  to  be  such  that, 
in  air,  the  number  striking  a  single  square  centimeter  of  surface 
per  second  would  fill  no  less  than  twenty  liters. 

Liquefaction  of  Gases:  Critical  Temperature.  —  Finally, 
gases  can  be  liquefied  by  sufficient  cooling  and  compression.  This 
fact  compels  us  to  suppose  that,  after  all,  even  gaseous  molecules 
have  a  tendency  to  cohesion.  This  cohesion  is  scarcely  perceptible 
so  long  as  the  gas  is  warm  and  is  diffuse.  Thus,  2  liters  of  oxygen  at 
1  atmosphere  pressure,  when  subjected  to  2  atmospheres  pressure, 
give  0.9991  liters  instead  of  1  liter.  The  additional  contraction 
of  0.0009  liters  (0.9  c.c.)  is  due  to  the  effect  of  cohesion  when 
the  molecules  are  thus  crowded  closer  together.  The  gases  which 
are  more  easily  liquefied  than  is  oxygen  show  greater  effects. 
Thus,  2  liters  of  sulphur  dioxide  at  760  mm.,  when  subjected  to  2 
atmospheres  pressure,  give  only  0.974  liters,  showing  a  contraction 
due  to  cohesion  of  26  c.c.  These  data  refer  to  0°.  At  lower 
temperatures  the  contractions  due  to  cohesion  become  rapidly 
greater.  This  cohesion  is  not  of  the  nature  of  gravitational 
attraction. 

We  can  readily  understand,  therefore,  that  when  the  kinetic 
energy  of  the  molecules  is  sufficiently  reduced  by  cooling  (namely, 
to,  or  below  the  critical  temperature,  see  below),  and  the  molecules 
are  brought  sufficiently  close  together,  the  tendency  of  the  mole- 
cules to  cohere  causes  the  gas  to  condense  and  assume  the  liquid 
form.  In  1869  Andrews  found  that  carbon  dioxide  could  be 
liquefied  at  0°  by  38  atmospheres  pressure,  and  at  30°  by  71  atmos- 
pheres, but  that  above  31.35°  it  could  not  be  liquefied  by  any 
pressure.  He  discovered  that  each  gas  has  a  critical  temperature, 
as  he  called  it.  For  carbon  dioxide,  this  temperature  can  be  ob- 
served by  placing  a  heavy-walled,  glass  tube  (Fig.  37),  half -filled 
with  liquid  carbon  dioxide,  in  a  beaker  of  water,  and  gradually 


STRUCTURE   AND   BEHAVIOR   OF   MATTER  79 

raising  the  temperature  of  the  latter.  At  31.35°,  the  surface 
between  the  liquid  and  gas  becomes  hazy  and  vanishes.  When 
the  temperature  falls  once  more,  the  surface  re-appears  at  31.35°. 
This  shows  that,  with  Faraday's  "permanent"  gases,  a  tempera- 
ture below  the  critical  point  had  not  been  employed. 

The  critical  temperature  of  oxygen  is  —118°,  of 
hydrogen  —234°,  of  carbon  dioxide  31.35°,  of  sulphur 
dioxide  156°,  of  water  358°. 

Another  Deviation  from  the  Laws  of  Gases. 
A  Perfect  Gas.  —  It  may  be  added  that  when  a  gas 
is  already  under  very  high  pressure,  and  very  closely 
packed,  an  increase  in  the  pressure  does  not  produce  quite 
as  great  a  diminution  in  volume  as  Boyle's  law  leads  us  to 
expect.  This  reminds  us  that  we  are  diminishing  only  FIG.  37. 
the  space  between  the  molecules,  and  not  the  volumes  of  the  mole- 
cules themselves,  and  therefore  not  the  total  volume  of  the  gas. 
When,  on  severe  compression,  the  volume  occupied  by  the  mole- 
cules themselves  has  become  an  appreciable  fraction  of  the  whole 
volume,  additional  compression  does  not  affect  the  whole  volume, 
and  the  contraction  is  smaller  than  Boyle's  law  would  indicate. 
Thus,  2  liters  of  hydrogen,  even  at  one  atmosphere  pressure,  when 
subjected  to  two  atmospheres  pressure,  give  1.0006  liters,  instead  of 
1  liter. 

The  last  two  effects  (namely,  those  due  to  the  tendency  to 
cohesion  of,  and  the  space  occupied  by  the  molecules)  are  called 
deviations  from  the  laws  of  gases.  In  consequence  of  these  individ- 
ual deviations,  there  are  not  exactly  equal  numbers  of  molecules  in 
equal  volumes  of  any  two  different  gases,  at  the  same  temperature 
and  pressure.  An  imaginary  gas,  which  exhibits  neither  deviation, 
called  a  perfect  gas,  is  often  referred  to  in  discussing  the  behavior 
of  gases. 

Summary.  —  We  may  now  summarize  the  principal  facts 
about  gases  in  mass,  appearing  in  italics  above,  with  the  corre- 
sponding features  of  the  molecular  relations,  in  heavy  type,  which 
we  have  added  one  by  one. 


80 


COLLEGE    CHEMISTRY 


Facts  About  Gases  in  Mass. 


Compressibility  . 

Diffusibility.  .  . 

Permeability  .  . 

Non-settling  .  . 

Homogeneity  .  . 

Pressure    .    .  .  . 

Boyle's  law  .  .  . 


Charles'  law. 


Above  and  other  facts 
Law  of  diffusion  .  . 
Gases  can  be  liquefied 


Corresponding  Relations  of  Molecules. 


Vacuum  +  molecules  widely  separated. 

Molecules  in  rapid  motion. 

Empty  space  relatively  large. 

Molecules  perfectly  elastic. 

Molecules  of  any  one  substance  closely  alike. 

Due  to  impacts  of  molecules. 

Pressure  proportional  to  concentration  of  the 

molecules.     Molecules  move  in  straight  lines 

and,  when  widely  scattered,  show  no  tendency 

to  cohesion  or  to  repulsion. 
Rise  in  temperature  increases  the  velocity,  and 

therefore  the  kinetic  energy  of  the  molecules. 
Avogadro's  law. 

Molecules  do  possess  a  tendency  to  cohesion, 
which  becomes  conspicuous  when  they  are 
cooled  and  closely  crowded  together. 


History  of  the  Kinetic  Molecular  Theory.  —  This  theory 
was  first  suggested  by  Daniel  Bernoulli  (1738),  who  explained  by 
its  means  the  pressure  and  compressibility  of  gases.  Lomonossov 
(1748)  developed  the  theory  very  completely  and  by  means  of  it 
explained  Boyle's  law  and  the  effects  of  changes  in  temperature. 
He  also  anticipated  from  the  theory  the  existence  of  the  second 
deviation  from  the  law  of  gases  (1749),  a  discovery  usually  cred- 
ited to  Dupre  (1864).  He  likewise  pointed  out  that  there  was 
no  limit  to  the  maximum  velocity  of  a  molecule,  and  therefore  no 
upper  limit  of  temperature,  but  that  there  must  be  a  lower  limit 
(the  absolute  zero)  at  which  the  molecules  would  be  at  rest  (1744). 
This  work  was  entirely  forgotten,  until  attention  was  called  to  it 
in  1904  by  Menschutkin. 

Similar  views  were  expressed  by  Waterston  (1845),  but  were 
still  so  much  ahead  of  the  time  that  the  committee  of  the  Royal 
Society  did  not  approve  the  paper  for  publication,  and  it  was  dis- 
covered in  the  archives  of  the  society,  long  afterwards,  by  Lord 
Rayleigh.  The  development. of  the  theory,  so  far  as  it  applies  to 
heat,  is  therefore  credited  to  Joule  (1855-60)  and,  in  respect  to  all 
properties  of  gases,  to  Kronig  (1856)  and  Clausius  (1857),  who 
knew  nothing  of  the  earlier  work. 


STRUCTURE   AND   BEHAVIOR   OF  MATTER  81 

Molecular  Relations  in  Liquids.  —  The  fact  that  even  great 
pressures  produce  little  diminution  in  the  volume  of  a  liquid  shows 
that  the  free  space,  present  in  gases,  is  absent  in  liquids.  The 
measured  effects  of  various  pressures  show,  for  example  in  the  case 
of  water,  that  to  reduce  the  volume  to  one-half  would  require,  not 
doubling  the  pressure  as  in  a  gas,  but  increasing  it  from  1  to  10,000 
atmospheres.  The  molecules  of  a  liquid  are  actually  in  contact 
with  one  another. 

The  phenomena  connected  with  surface  tension,  such  as  co- 
herence into  drops,  show  that  cohesion  plays  a  much  larger  part 
in  liquids  than  in  gases.  On  the  other  hand,  liquids  which  are 
capable  of  mixing  (e.g.,  alcohol  and  water),  when  placed  above  one 
another  in  the  same  vessel,  do  mix,  slowly,  by  diffusion.  This 
indicates  that  motion  of  the  molecules,  although  much  impeded 
by  friction,  has  not  been  annihilated  by  cohesion.  The  escape  of 
vapor  —  that  is,  of  part  of  the  liquid  in  gaseous  form  —  likewise 
proves  that  the  molecules  in  the  liquid  are  in  motion.  The  rela- 
tions of  liquid  and  vapor  can  be  discussed  most  effectively  in  the 
next  chapter,  in  connection  with  the  case  of  water  and  steam. 

Molecular  Relations  in  Solids.  —  The  properties  of  solids 
differ  from  those  of  liquids  chiefly  in  the  fact  that  the  solid  has 
a  definite  form  of  which  it  can  be  deprived  only  with  difficulty. 
This  we  may  explain  in  accordance  with  the  kinetic  hypothesis  by 
the  supposition  that  the  cohesion  in  solids  is  very  much  more 
prominent  than  in  liquids.  We  obtain  solids  from  liquids  by 
cooling  them;  in  other  words,  by  diminishing  the  kinetic  energy 
and  therefore  the  velocity  of  the  particles.  The  cohesive  tendency 
of  the  latter  is  thus  able  to  make  itself  felt  to  a  greater  extent.  If, 
conversely,  we  heat  a  solid,  or,  according  to  the  hypothesis,  if  we 
increase  the  speed  with  which  the  particles  move,  the  body  first 
melts  and  gives  a  liquid,  and  this  finally  boils  and  becomes  a  gas. 
The  intrinsic  cohesion  of  the  particular  substance  can  undergo  no 
change,  but  the  increasing  kinetic  energy  of  the  particles  steadily 
and  continuously  obliterates  its  effects.  Yet  some  motion  still 
survives  in  a  solid.  Thus  we  find  that  when  the  layer  of  silver  is 
stripped  from  a  very  old  piece  of  electroplate,  the  presence  of  this 
metal  in  the  German  silver  or  copper  basis  of  the  article  is  easily 
demonstrated. 


82  COLLEGE    CHEMISTRY 

The  tendency  of  all  solids  to  assume  crystalline  forms,  which 
show  definite  cleavage  and  other  evidences  of  structure,  distin- 
guishes them  sharply  from  liquids.  The  force  of  cohesion  in 
liquids  is  exercised  equally  in  different  directions.  In  solids  it 
must  differ  in  different  directions  in  order  that  structure  may  re- 
sult. Since  each  substance  shows  an  individual  structure  of  its 
own,  these  directive  forces  must  have  special  values  in  magnitude 
and  direction  in  each  substance. 

Crystallization.  —  A  crystal  arises  by  growth.  When  the 
process  is  watched,  as  it  occurs  in  a  melted  solid  or  an  evaporating 
solution,  the  slow  and  systematic  addition  of  the  material  in  lines 
and  layers,  as  if  according  to  a  regular  design,  is  one  of  the  most 
beautiful  and  interesting  of  natural  phenomena.  The  fern-like 
patterns  produced  by  ice  on  a  window-pane  show  the  general 
appearance  characteristic  of  crystallization  in  a  thin  layer.  A 
larger  mass  in  a  deep  vessel  gives  forms  which  are  geometrically 
more  perfect.  From  its  very  incipiency  the  crystal  has  the  same 
form  as  when,  later,  its  outlines  can  be  distinguished  by  the  eye. 
Hence  the  outward  form  is  only  an  expression  of  a  specific  internal 
structure  which  the  continual  reproduction  of  the  same  outward 
form  on  a  larger  and  larger  scale  leaves  as  a  memorial  of  itself  in 
the  interior. 

Crystal  Forms.  —  Crystalline  form  is  continually  used  in 
identifying  (pp.  2,  12,  19)  the  substances  produced  in  chemical 
actions.  The  classification  of  crystalline  forms  is  carried  out 
according  to  the  degree  of  symmetry  of  the  crystals: 

1.  Regular  system.  5.   Monosymmetric,  or 

2.  Square  prismatic  system.  monoclinic  system. 

3.  Hexagonal  system.  6.   Asymmetric,  or 

4.  Rhombic  system.  triclinic  system. 

The  regular  system  presents  the  most  symmetrical  figures  of  all. 
Some  forms  which  commonly  occur  are  the  octahedron  (Fig.  38) 
shown  by  alum,  the  cube  (Fig.  39)  affected  by  common  salt,  and 
the  dodecahedron  (Fig.  40)  frequently  assumed  by  the  garnet. 

The  square  prismatic  system  includes  less  symmetrical  forms  than 
the  previous  one,  since  the  crystals  are  lengthened  in  one  direction. 


STRUCTURE   AND   BEHAVIOR   OF   MATTER 


83 


Fig.  41  shows  the  condition  in  which  zircon  (ZrSi04),  which  fur- 
nishes us  with  the  basis  of  certain  incandescent  illuminating 
arrangements,  occurs  in  nature.  The  form  of  ordinary  hydrated 
nickel  sulphate  (NiSO4,6H20)  is  similar  to  this. 


FIG.  38. 


Fia.  40. 


FIG.  41. 


The  hexagonal  system,  like  the  preceding,  frequently  exhibits 
elongated  prismatic  forms,  but  the  section  of  the  crystals  is  a  hexa- 
gon, instead  of  a  square,  and  the  termination  is  a  six-sided  pyramid. 
Quartz  (Fig.  42),  or  rock  crystal,  is  the  most  familiar  mineral  in 


FIG.  42.         FIG.  43. 


FIG.  44. 


FIG.  45. 


this  system.  Calcite  (CaCO3),  which  is  chemically  identical  with 
chalk,  or  marble,  takes  forms  known  as  the  scalenohedron  (Fig.  43) 
and  rhombohedron  (Fig.  44),  which  are  classified  in  a  subdivision 
of  this  system.  Indeed,  recently  it  has  become  common  to  erec* 


FIG.  46. 


FIG.  47. 


FIG.  48. 


this  into  a  separate  system  (the  trigonal),  in  which  both  quartz  and 
calcite  are  included. 

The  rhombic  system  includes  the  natural  forms  of  the  topaz, 
and  of  sulphur  (Fig.  7,  p.  12),  as  well  as  that  of  potassium  perman- 
ganate (Fig.  45),  potassium  nitrate  (Fig.  46),  and  many  other 


84  COLLEGE    CHEMISTRY 

substances.  These  crystals  exhibit  a  good  deal  of  symmetry,  but 
their  section  is  always  rhombic,  and  hence  the  name. 

The  monosymmetric  system  exhibits  forms  which  have  but  one 
plane  of  symmetry.  Gypsum  (Fig.  47),  which  is  hydrated  calcium 
sulphate  CaS04,2H20,  and  feldspar  (Fig.  3,  p.  2)  are  minerals  pos- 
sessing forms  of  this  kind.  Tartaric  acid,  rock  candy  (Fig.  48), 
potassium  chlorate,  and  hydrated  sodium  carbonate  (washing 
soda)  belong  to  this  system. 

The  asymmetric  system  includes  forms  which  have  no  plane  of 
symmetry  whatever.  Blue  vitriol  (Fig.  52,  p.  95),  CuS04,5H20,  is 
one  of  the  most  familiar  substances  of  this  kind. 

Exercises.  —  The  text  cannot  be  understood  unless  some 
problems  involving  the  laws  of  gases  are  actually  worked. 

1.  Reduce  189  c.c.  of  gas  at  15°  and  750  mm.  to  0°  and  760  mm. 

2.  Reduce  110  c.c.  of  gas  at  -  5°  and  741  mm.  to  0°  and  760  mm. 

3.  Convert  500  c.c.  of  gas  at  25°  and  700  mm.  to  18°  and  745  mm. 

4.  Reduce  250  c.c.  of  gas  (standing  over  water)  at  22°  and  755 
mm.  to  the  dry  condition  and  to  0°  and  760  mm. 

5.  The  density  of  a  substance  referred  to  air  is  3.2.     What  is  the 
density  referred  to  hydrogen?     What  will  be  the  volume  occupied 
by  10  g.  of  the  substance  at  20°  and  752  mm.? 

6.  Describe  two  ways  of  obtaining  crystals  of  a  substance. 


CHAPTER  VIII 
WATER 

WATER  is  as  necessary  to  life  as  is  oxygen.  The  human  body 
is  saturated  with  it  and,  to  make  up  for  evaporation,  as  well  as  to 
aid  in  digestion  and  other  life  processes,  i$  is  a  necessary  part  of 
our  food.  The  ocean  covers  three-fourths  of  the  earth's  surface, 
and  the  "dry"  land  is,  fortunately,  far  from  being  really  dry. 

Physical  Properties  of  Water.  —  A  deep  layer  of  water  has  a 
blue  or  greenish-blue  color.  At  a  pressure  of  760  mm.,  water  ex- 
ists as  a  liquid  between  0°  and  100°.  Below  0°  it  becomes  solid, 
above  100°  a  gas.  Of  all  chemical  substances  it  is  the  one  which  we 
use  most,  so  that  its  physical  properties,  discussed  below,  should 
be  studied  attentively.  Then,  too,  what  is  said  of  water  is  in 
general  true  of  all  other  liquids,  from  which  it  differs  only  in 
details. 

The  quantity  of  heat  required  to  raise  one  gram  of  water  one 
degree  in  temperature,  at  15°,  is  called  a  calorie,  the  unit  quantity 
of  heat.  The  specific  heat  of  any  substance  being  the  quantity 
of  heat  required  to  raise  the  temperature  of  one  gram  of  the  sub- 
stance one  degree,  the  specific  heat  of  water  is  1.  The  values  for 
other  substances  are  all  smaller  (e.g.,  limestone  0.2).  Thus  the 
temperature  of  large  masses  of  water,  such  as  lakes  and  seas, 
changes  more  slowly,  and  within  a  smaller  range,  than  that  of  the 
rocks  and  soil  composing  the  land.  The  more  constant  tempera- 
ture of  the  water  tends  to  regulate  that  of  the  air,  and  hence  the 
climate  of  an  island  is  less  variable  from  season  to  season  than  is 
that  of  a  continent. 

Ice.  —  The  raising  or  lowering  of  the  temperature  of  a  gram  of 
water  through  one  degree  involves  the  addition  or  removal  of  one 
calorie  of  heat.  The  conversion,  however,  of  a  gram  of  water  at 
0°  to  a  gram  of  ice  at  0°  requires  the  removal  of  79  calories.  The 

85 


86  COLLEGE    CHEMISTRY 

mere  melting  of  a  gram  of  ice  causes  an  absorption  of  heat  to  the 
same  amount,  called  the  heat  of  fusion  of  ice.  At  0°  a  mixture  of 
ice  and  water  will  remain  in  unchanged  proportions  indefinitely. 
Any  cause  which  tends  permanently  to  lower  or  raise  the  tempera- 
ture by  a  fraction  of  a  degree,  however,  will  bring  about  the  disap- 
pearance of  the  water  or  of  the  ice,  respectively.  This  temperature 
is  called  the  melting-  or  the  freezing-point. 

Water  can  be  cooled  below  0°  (supercooled)  without  beginning 
to  freeze,  unless  it  is  stirred,  or  "inoculated"  by  the  addition  of  a 
piece  of  ice.  Hence,  the  freezing-point  is  not  defined  as  the  point 
at  which  ice  begins  to  form,  for  that  point  varies,  and  is  always 
below  0°,  but  as  the  temperature  of  a  mixture  of  ice  and  water. 

Steam.  —  At  atmospheric  pressure,  water  passes  into  steam 
rapidly  at  100°,  but  at  lower  temperatures,  and  even  when  frozen, 
it  does  the  same  thing  more  slowly.  It  changes  into  steam,  how- 
ever, only  when  the  necessary  supply  of  heat  is  forthcoming. 
One  gram  of  water  at  100°,  in  turning  into  a  gram  of  steam  at  100°, 
takes  up  540  calories.  This  is  called  its  heat  of  vaporization. 
Steam,  in  fact,  contains  much  more  internal  energy  than  an  equal 
weight  of  water  at  the  same  temperature,  just  as  water,  in  turn, 
contains  more  energy  than  ice. 

Steam  is  a  colorless,  invisible  gas.  The  visible  cloud  of  fog, 
seen  when  steam  escapes  into  cold  air,  is  composed  of  minute 
drops  of  water,  formed  by  condensation,  and  visible  because  they 
have  surfaces  and  reflect  light. 

The  States  of  Matter:  Transition  Points.  —  Most  sub- 
stances are  known  in  three  different  states  of  aggregation,  solid 
(crystalline),  liquid,  and  gaseous.  There  is  no  magic  about  the 
number,  three,  however.  Thus,  sulphur  has  a  vapor  state,  two 
liquid  states,  and  several  solid  forms.  There  are  even  five  forms 
of  ice,  and  most  solids  probably  exist  in  several  different  states. 

All  transitions  from  one  state  to  another  take  place  at  some 
definite  temperature  (when  the  pressure  is  fixed).  Such  temper- 
atures, when  referring  to  the  change  from  the  solid  to  the  liquid, 
and  from  the  liquid  to  the  gaseous  state  are  called  the  melting- 
point  or  freezing-point,  and  the  boiling-point,  respectively,  or  in 
general,  are  known  as  transition  points. 


WATER 


87 


Aqueous  Tension  and  Vapor  Pressure.  —  The  quantity 
of  the  vapor  present  is  defined  by  the  gaseous  pressure  it  exercises, 
the  value  being  called  the  vapor  pressure  of  water  vapor  (or  of 
the  vapor  of  any  other  volatile  substance)  in  the  location  in 
question. 

The  most  significant  fact  about  vapor  pressure  is  that,  when  ex- 
cess of  the  liquid  is  present,  the  pressure  of  the  vapor  quickly  reaches 
a  definite  maximum  value  for  each  temperature.  In  the  absence 
of  excess  of  the  water,  less  than  this  maximum  pressure  may  exist. 
More  than  the  maximum  pressure  proper  to  a  given  temperature,  if 
produced  by  compression,  cannot  be  maintained,  however,  for  a 
part  of  the  vapor  condenses  to  the  liquid  state.  The  magnitude 
of  this  maximum  vapor  pressure,  at  a  given  temperature,  depends 
on  the  ability  of  the  particular  liquid  to  generate  vapor.  This  max- 
imum vapor  pressure  is  held,  therefore,  to  represent  the  vapor 
tension  of  the  liquid,  at  the  given  temperature,  and  this  is  a  specific 
property  of  the  substance. 

The  vapor  tension  may  be  shown  by  allowing  a  few  drops  of 
water  to  ascend  into  a  barometric  vacuum  (Fig.  49).  The  tube 
on  the  left  shows  the  mercury  when  nothing  presses 
on  its  surface.  The  tube  on  the  right  shows  the 
result  of  admitting  the  water.  The  difference  in 
the  height  of  the  two  columns  gives  the  value  of  the 
vapor  pressure  of  the  water  vapor.  With  excess  of 
water,  the  value  is  that  of  the  vapor  tension,  called, 
in  the  case  of  water,  the  aqueous  tension. 

The  jacket  surrounding  the  tube  on  the  right 
enables  us,  by  adding  ice  or  warm  water,  to  main- 
tain any  temperature  between  0°  and  100°.  When 
ice  is  used  outside,  and  a  piece  of  it  is  introduced 
into  the  vacuum,  the  vapor  it  gives  off  quickly 
reaches  a  pressure  of  4.5  mm.  The  vapor  pressure 
of  the  ice  takes  the  place  of  4.5  mm.  of  mercury  in 
balancing  the  atmospheric  pressure,  and  so  the  mer- 
cury column  falls  by  this  amount.  Similarly,  water 
at  10°  causes  a  fall  of  9.1  mm.  and  at  20°  of  17.4  mm.,  so  that  these 
represent  the  mercury-height  values  of  the  aqueous  tension  at 
these  temperatures.  The  quantity  of  water  used  makes  no  differ- 
ence, so  long  as  a  little  more  is  present  than  is  required  to  fill  the 


FIG.  49. 


88  COLLEGE   CHEMISTRY 

available  space  with  vapor.  With  ether,  instead  of  water,  at 
10°,  the  fall  is  28.7  mm. 

With  water  at  higher  temperatures  the  fall  of  the  mercury  col- 
umn becomes  much  greater.  At  50°  it  is  92  mm.,  at  70°  it  is  233.3 
mm.,  at  90°  it  is  525.5  mm.,  and  at  100°  it  is  760  mm.,  or  one  at- 
mosphere. At  121°  the  aqueous  tension  is  two  atmospheres,  at 
180°  it  is  ten  atmospheres  (see  Appendix  IV). 

When  water  at  a  certain  temperature  has  given  the  full  amount  of 
water  vapor  to  the  space  above  it  that  its  aqueous  tension  permits, 
we  say  that  the  space  is  saturated  with  vapor.  That  concentration 
of  vapor  which  constitutes  saturation  varies  with  the  temperature 
of  the  water  and  depends,  therefore,  solely  on  the  power  of  the  water 
to  give  off  vapor.  It  has  nothing  to  do  with  the  size  of  the  space, 
and  is  even  independent  of  other  gases  the  space  may  already  con- 
tain. Thus,  if  a  little  air  is  first  placed  above  the  dry  mercury 
(Fig.  49),  causing  it  to  fall,  the  additional  depression  produced  by 
adding  water  is  the  same  as  if  the  air  had  been  absent  (p.  72.  See 
footnote  to  p.  11). 

Water  Vapor  in  the  Air.  —  The  space  immediately  above  the 
surface  of  the  ground,  which  is  mainly  occupied  by  atmospheric 
air,  is,  on  an  average,  less  than  two-thirds  saturated  with  water 
vapor.  That  is  to  say,  such  air,  when  enclosed  in  a  vessel  con- 
taining water,  will  take  up  about  one-half  more  than  it  already 
contains.  The  vapor  of  water  at  100°  in  an  open  vessel  displaces 
the  air  entirely  and,  if  the  required  heat  is  furnished,  the  liquid 
boils. 

All  our  substances  and  apparatus  have  traces  of  water,  derived 
from  the  atmosphere,  condensed  on  their  surfaces.  This  water  is, 
in  a  sense,  in  an  abnormal  condition,  for  it  does  not  evaporate  even 
in  dry  air.  It  is  observed  to  pass  off  in  vapor,  however,  when  we 
have  occasion  to  heat  the  substance  or  apparatus. 

Molecular  Relations  of  Liquid  and  Vapor.  —  When  the 
water  was  introduced  above  the  barometric  column,  the  vapor,  or 
gaseous  water,  could  have  resulted  only  from  the  spontaneous 
motion  of  the  molecules  in  the  liquid.  Some  of  the  molecules, 
moving  near  the  surface,  went  off  into  the  space  above  the  water 
and  became  gaseous.  To  be  consistent,  we  must  also  conclude 


WATER  89 

that  the  vapor  above  the  water  is  not  composed  of  the  same  set  of 
molecules  one  minute  as  it  was  during  the  preceding  minute. 
Their  motions  must  cause  many  of  them  to  plunge  into  the  liquid, 
while  others  emerge  and  take  their  places.  When  the  water  is  first 
introduced,  there  are  no  molecules  of  vapor  in  the  space  at  all,  so 
that  emission  from  the  water  predominates.  The  pressure  of  the 
vapor  increases  as  the  concentration  of  the  molecules  of  vapor 
becomes  greater,  hence  the  mercury  column  falls  steadily.  At  the 
same  time  the  number  of  gaseous  molecules  plunging  into  the 
water  per  second  must  increase  in  proportion  to  the  degree  to 
which  they  are  crowded  in  the  vapor.  The  rate  at  which  mole- 
cules return  to  the  water  thus  begins  at  zero,  and  increases  steadily; 
the  rate  at  which  molecules  leave  the  water  maintains  a  constant 
value.  Hence  the  rate  at  which  vapor  molecules  enter  the  water 
must  eventually  equal  that  at  which  other  water  molecules  leave 
the  liquid.  At  this  point,  occasion  for  visible  changes  ceases  and 
the  mercury  comes  to  rest.  We  are  bound  to  think,  however, 
of  the  exchange  as  still  going  on,  since  nothing  has  occurred  to 
stop  it.  The  condition  is  not  one  of  rest  but  of  rapid  and  equal 
exchange.  Such,  described  in  terms  of  molecules,  is  the  state  of 
affairs  which  is  characteristic  of  a  condition  of  equilibrium.  The 
condition  is  dynamic,  and  not  static. 

Equilibrium.  —  This  term  is  used  so  often  in  chemistry,  and  is 
used  in  so  unfamiliar  a  sense,  that  the  reader  should  consider 
attentively  what  it  implies.  Three  things  are  characteristic  of  a 
state  of  equilibrium: 

1.  There  are  always  two  opposing  tendencies  which,  when  equi- 
librium is  reached,  balance  each  other.  In  the  foregoing  instance, 
one  of  these  is  the  hail  of  molecules  leaving  the  liquid,  which  is 
constant  throughout  the  experiment.  It  represents  the  vapor  ten- 
sion of  the  liquid.  The  other  is  the  hail  of  returning  molecules, 
which,  at  first,  increases  steadily  as  the  concentration  of  the  vapor 
becomes  greater.  This  is  the  vapor  pressure  of  the  vapor.  These 
have  the  effect  of  opposing  pressures  and,  when  the  latter  becomes 
equal  to  the  former,  equilibrium  is  established.  In  all  cases  of 
equilibrium  we  shall  symbolize  the  two  opposing  tendencies  by  two 
arrows,  thus: 

Water  (liq.)  ^  Water  (vapor). 


90  COLLEGE   CHEMISTRY 

2.  Although  their  effects  thus  neutralize  each  other  at  equilib- 
rium, both  tendencies  are  still  in  full  operation.  In  the  case  in  point, 
the  opposing  hails  of  molecules  are  still  at  work,  but  neither  can 
effect  any  visible  change  in  the  system.  Equilibrium  is  a  state, 
not  of  rest,  but  of  balanced  activities. 

3  (and  this  is  the  chief  mark  of  equilibrium).  A  slight  change  in 
the  conditions  produces,  never  a  great  or  sharp  change,  but  always, 
and  instantly,  a  corresponding  small  change  in  the  state  of  the 
system.  The  change  in  the  conditions  accomplishes  this  by  favoring 
or  disfavoring  one  of  the  two  opposing  tendencies.  Thus,  for  ex- 
ample, when  the  temperature  of  a  liquid  is  raised,  the  kinetic  energy 
of  its  molecules  is  increased,  the  rate  at  which  they  leave  its  surface 
becomes  greater,  the  vapor  tension  increases  and,  hence,  a  greater 
concentration  of  vapor  can  be  maintained.  The  system,  therefore, 
quickly  reaches  a  new  state  of  equilibrium  in  which  a  higher  vapor 
pressure  exists.  A  heap  of  matter  on  a  table  is  not  in  equilib- 
rium, because  addition  of  more  material  produces  no  response 
until,  when  a  very  great  quantity  is  added,  the  table  breaks. 
But  a  body  on  the  scales  is  in  equilibrium,  for  the  addition  of 
the  smallest  particle  produces  a  corresponding  inclination  of  the 
beam. 

In  the  preceding  illustration,  the  evaporating  tendency  was 
favored  by  a  rise  in  temperature.  As  an  example  of  a  change  in 
conditions  disfavoring  one  tendency,  take  the  case  where  the  liquid 
is  placed  in  an  open,  shallow  vessel.  Here  the  condensing  tendency 
is  markedly  discouraged,  for  there  is  practically  no  return  of  the 
emitted  molecules.  Hence  complete  evaporation  takes  place.  Ele- 
vation of  the  temperature  hastens  the  process.  A  draft  insures  the 
total  prevention  of  all  returns,  and  has  therefore  the  same  effect. 
The  two  methods  of  assisting  the  displacement  of  an  equilibrium, 
and  particularly  the  second,  in  which  the  opposed  process  is  weak- 
ened and  the  forward  process  triumphs  solely  on  this  account, 
should  be  noted  carefully.  They  are  applied  with  surprising  effec- 
tiveness in  the  explanation  of  chemical  phenomena  (see  Chaps.  XIV 
and  XVIII). 

Water  as  a  Solvent.  —  One  of  those  physical  properties  of 
water  which  are  most  used  in  chemical  work  is  its  tendency  to  dis- 
solve many  substances.  This  subject  is  so  important  and  exten- 


WATER  91 

sive  that  we  shall  presently  devote  a  complete  chapter  to  some  of 
its  simpler  and  more  familiar  aspects. 

Natural  Waters.  —  The  foreign  material  in  natural  waters  is 
divided  into  dissolved  matter  and  suspended  matter.  Rain-water, 
collected  after  most  of  the  dust  has  been  carried  down,  is  the 
purest  natural  water.  It  contains,  however,  nitrogen,  oxygen,  and 
carbon  dioxide  dissolved  from  the  air.  Sea-water  holds  about  3.6 
per  cent  of  dissolved  material.  River  and,  especially,  well  waters 
dissolve  various  substances  during  their  progress  over  or  under 
the  surface  of  the  ground.  They  often  contain  calcium  sulphate, 
calcium  bicarbonate,  and  compounds  of  magnesium,  and  are  then 
described  as  hard.  Sometimes  they  contain  compounds  of  iron, 
and  sometimes  they  are  effervescent  and  give  off  carbon  dioxide. 
These  are  called  mineral  waters. 

Many  river  waters  contain  large  amounts  of  clay  and  organic 
matter  (often  due  to  admixture  of  sewage)  suspended  in  them. 
It  is  not  the  organic  matter  which  is  deleterious,  but  the  bacteria 
of  putrefaction  and  disease  which  are  present  also,  and  are  usually 
for  the  most  part  attached  to  the  particles  of  suspended  matter. 
Cholera  and  typhoid  fever  are  often  spread  by  the  drinking  of 
water  into  which  sewage,  infected  by  other  patients  suffering  from 
these  diseases,  has  been  allowed  to  enter.  Clay  can  be  seen,  and 
renders  the  water  turbid,  but  organic  matter  and  bacteria  may  be 
present  in  water  which  looks  perfectly  clear. 

Purification  from  Suspended  Matter.  —  The  suspended 
impurities  may  be  removed  by  filtration.  On  a  large  scale,  beds  of 
gravel  are  employed,  but  this  treatment  will  not  remove  all  bac- 
teria. In  many  cases  small  amounts  of  alum,  or  alum  and  lime, 
or  ferrous  sulphate  (copperas)  and  lime,  are  added.  These  pro- 
duce slimy  precipitates  which  assist  in  the  elimination  of  fine,  sus- 
pended inorganic  and  organic  matter,  including  practically  all  the 
bacteria.  This  is  called  the  coagulation  treatment  (q.v) .  The  whole 
suspended  matter  is  then  allowed  to  settle,  which  it  does  very 
quickly,  in  large  reservoirs.  The  remaining  organisms  may  be 
destroyed  by  adding  a  little  bleaching  powder  (q.v.),  before  the 
water  is  distributed.  Ozone  and  ultra-violet  light  are  used  for 
the  same  purpose. 


92 


COLLEGE   CHEMISTRY 


In  the  household,  the  Pasteur  filter  is  the  most  compact  and 
efficient  appliance.  The  water  enters  at  the  top  (Fig.  50),  and  is 
forced  inwards  by  its  own  pressure  through  the  pores  of  a  cylinder 
of  unglazed  porcelain.  The  cylinder  must  be  taken 
out,  and  its  exterior  cleaned  daily  with  a  brush,  to 
remove  the  mud  and  organisms  which  collect  on  its 
outer  surface.  If  this  is  not  done,  the  organisms 
multiply  and  soon  the  filter  pollutes  the  water  instead 
of  purifying  it. 

Most  organisms  can  be  killed  by  boiling  the  un- 
filtered  water  for  10  or  15  minutes,  although  a  second 
boiling  is  needed  in  the  case  of  some. 

Purification  from  Dissolved  Matter.  —  Filtra- 
tion does  not  remove  dissolved  matter,  and  therefore 
does  not  soften  hard  water  (q.v.). 

Pure  water  for  chemical  purposes  is  prepared  by 
distillation  and,  in  fact,  liquids  other  than  water  are 
usually  purified  by  the  same  process  (Fig.  51).  The 
steam  is  condensed  by  cold  water  circulating  in  the 
jacket,  and  contains  only  gases  dissolved  from  the 
air.  Dissolved  solids  remain  in  the  flask.  Distilled  water  quickly 
dissolves  traces  of  glass  or  porcelain,  so  that  the  purest  water  is 
obtained  by  using  quartz  or  platinum  for  the  condenser  tube  and 
receiving  vessel.  Tin  is  the  best  of  the  less  expensive  materials. 

Chemical  Properties  of  Water.  —  Water  is  so  very  frequently 
used  in  chemical  experiments  in  which  it  is  a  mere  mechanical  ad- 
junct, that  the  beginner  has  difficulty  in  distinguishing  the  cases  in 
which  it  has  itself  taken  part  in  the  chemical  interaction.  The 
four  kinds  of  chemical  activity  which  it  shows  should  therefore 
receive  careful  notice: 

1.  Water  is  a  relatively  stable  substance. 

2.  It  combines  with  many  oxides,  forming  bases  or  acids. 

3.  It  combines  with  many  substances,  chiefly  salts,  forming 
hydrates. 

4.  It  interacts  with  some  substances  in  a  way  described  as  hy- 
drolysis.    This  property  will  not  be  discussed  until  a  characteristic 
case  is  encountered. 


FIG.  50. 


WATER 


93 


Water  a  Stable  Compound:  Dissociation.  —  In  the  case  of 
a  compound,  the  first  chemical  property  to  be  given  is  always, 
whether  the  substance  is  relatively  stable  or  unstable.  Usually  the 


FIG.  51. 

specification  is  in  terms  of  the  temperature  required  to  produce 
noticeable  decomposition.  Thus,  potassium  chlorate  gives  off 
oxygen  at  a  low  red  heat.  Now,  water  vapor,  when  heated,  is 
progressively  decomposed  into  hydrogen  and  oxygen,  yet  at  2000° 
the  decomposition  reaches  only  1.8  per  cent,  and  reunion  occurs  as 
the  temperature  is  lowered.  The  two  arrows  in  the  equation  indi- 
cate that  the  action  may  proceed  in  either  direction  —  is  reversible : 

H20<=±2H  +  0. 

A  decomposition  which  thus  proceeds  at  higher  temperatures, 
while  at  lower  temperatures  combination  of  the  constituents  can 
take  place,  is  called  a  dissociation.  The  decomposition  of  potas- 
sium chlorate  (p.  27)  is  not  a  dissociation  because  it  is  not  revers- 
ible; oxygen  gas  will  not  under  any  known  circumstances  unite  with 
potassium  chloride. 


94  COLLEGE   CHEMISTRY 

Union  of  Water  with  Oxides.  —  1.   Sodium  oxide   (Na^O) 
unites  violently  with  water  to  form  sodium  hydroxide: 


The  slaking  of  quicklime  is  a  more  familiar  action  of  the  same  kind: 

H2O^Ca(OH)2. 


No  other  products  are  formed.  The  clouds  of  condensing  steam 
produced  in  the  second  instance  are  due  to  evaporation  of  a  part 
of  the  water  by  the  heat  produced  in  the  formation  of  calcium 
hydroxide.  The  aqueous  solutions  of  these  two  products  have  a 
soapy  feeling,  and  turn  red  litmus  (a  vegetable  extract)  blue,  and 
the  substances  therefore  belong  to  the  class  of  alkalies  or  bases. 
Very  many  hydroxides,  which  are  of  the  same  nature,  for  example 
ferric  hydroxide  Fe(OH)3  and  tin  hydroxide  Sn(OH)2,  are  formed 
so  slowly  by  direct  union  of  the  oxide  and  water  that  they  are 
always  prepared  in  other  ways.  The  oxides  which,  with  water, 
form  bases  are  called  basic  oxides. 

2.  Some  oxides,  although  they  unite  with  water,  give  acids, 
which  are  products  of  an  entirely  different  character.  Phosphorus 
pentoxide  (p.  32)  and  sulphur  dioxide  are  of  this  class  and  yield 
phosphoric  acid  and  sulphurous  acid.  Such  oxides  are  commonly 
called  the  anhydrides  (Gk.,  without  water)  of  their  respective  acids. 
They  are  called  also  acidic  oxides: 

3H20-+2H3P04. 


The  acids  are  sour  in  taste  and  turn  blue  litmus  red. 

These  two  classes  of  final  products  are  so  different  that  we  make 
the  distinction  the  basis  for  classification  of  the  elements  present 
in  the  original  oxides.  The  elements,  like  sodium  and  iron,  whose 
oxides  give  bases,  are  called  metallic  elements;  those,  like  phos- 
phorus, whose  oxides  give  acids,  are  called  non-metallic  elements. 
The  distinguishing  words  are  selected  because  the  division  corre- 
sponds, in  a  general  way  at  least,  with  the  separation  into  two  sets 
to  which  merely  physical  examination  of  the  elementary  substances 
would  lead. 


WATER  95 

Hydrates.  —  Many  substances  when  dissolved  in  water  and 
recovered  by  spontaneous  evaporation  of  the  solvent  enter  into 
combination  with  the  liquid.  The  products,  which  are  solids,  are 
called  hydrates.  That  they  are  regular  chemical  compounds  is 
shown  by  the  following  two  facts:  (1)  These  compounds  show 
definite  chemical  composition  expressible  by  formulae  in  terms  of 
chemical  unit  weights  (atomic  weights)  of  the  constituent  ele- 
ments. The  proportions  in  solutions  and  other  physical  aggre- 
gations, except  by  chance,  cannot  be  expressed  by  means  of 
formulae.  (2)  A  hydrate  has  physical  properties  entirely  different 
from  those  of  the  water  (or  icejr  and  the 
other  substance  used  in  preparing  it.  It  is 
a  typical  compound,  formed  by  the  first 
variety  of  chemical  change  (p.  7).  Thus, 
cupric  sulphate,  often  called  anhydrous 
cupric  sulphate  to  distinguish  it  from  the 
compound  with  water,  is  a  white  substance 
crystallizing  in  shining,  colorless,  needle- 
like  prisms.  The  pentahydrate  (blue-stone 
or  blue  vitriol)  which  crystallizes  from  the 

aqueous  solution,  is  blue  in  color,  and  forms  larger  but  much  less 
symmetrical  (asymmetric  or  triclinic)  crystals  (Fig.  52) : 

CuSO4  +  5H20  <=»  CuS04,5H20. 

The  chemical  properties  show  hydrates  to  be  relatively  unstable. 
When  heated,  the  hydrates,  as  a  rule,  lose  none  of  the  constituents 
of  the  original  compound,  but  only  the  water,  in  the  form  of  vapor. 
When  melted,  or  when  dissolved  in  water,  the  hydrates  are  disso- 
ciated (p.  93)  into  water  and  the  original  substance.  The  aqueous 
solutions  made  from  the  anhydrous  substances  and  from  the  hy- 
drates have  identical  physical  and  chemical  properties.  Hence 
the  cheaper  of  the  two  forms  is  generally  purchased,  and  many 
of  the  chemicals  used  in  the  laboratory  are  in  the  form  of 
hydrates. 

In  consequence  of  the  ease  with  which  hydrates  give  up  water 
we  write  their  formulae  (e.g.,  CuS04,5H20)  so  that  the  water  and 
original  substance  are  separate.  A  formula  thus  modified,  so  as 
to  show  some  favorite  mode  of  behavior  of  the  substance,  is  called 
a  reaction  formula.  The  formula  Hi0CuS09,  which  would  show  the 


96  COLLEGE    CHEMISTRY 

same  proportions  by  weight,  is  never  employed,  because  its  use 
would  disguise  the  relation  of  the  substance  to  cupric  sulphate. 

The  Dissociation  of  Hydrates.  Efflorescence.  —  The  less 
stable  hydrates  dissociate  very  readily.  Thus  the  decahydrate 
of  sodium  sulphate,  Na2S04,10H2O  (Glauber's  salt),  loses  all  the 
water  it  contains  (effloresces)  when  simply  kept  in  an  open  vessel. 
When  kept  in  a  closed  bottle,  a  very  little  of  it  loses  water,  and 
then  the  decomposition  ceases.  The  cause  of  this  we  discover 
when  a  crystal  of  the  hydrate  is  placed  above  mercury,  like  the  ice 
or  water  in  Fig.  49  (p.  87).  It  shows  an  aqueous  tension  which  we 
can  measure.  At  9°  the  value  of  this  is  5.5  mm.  As  its  tempera- 
ture is  raised,  the  tension  increases.  When  the  temperature  is 
lowered,  on  the  other  hand,  the  tension  diminishes,  the  mercury 
rises,  and  a  part  of  the  water  enters  into  combination  again. 
Different  hydrates  show  different  aqueous  tensions  at  the  same 
temperature.  For  example,  at  30°,  that  of  water  itself  is  31.5  mm., 
strontium  chloride  SrCl2,6H20,  11.5  mm.;  cupric  sulphate  CuSO4, 
5H20,  12.5  mm.;  barium  chloride  BaCl2,2H2O,  4  mm. 

In  view  of  these  facts,  we  perceive  that  loss  of  water  by  efflores- 
cence is  like  evaporation,  excepting  that  it  is  a  chemical  decompo- 
sition and  not  a  physical  process.  Those  hydrates  which,  like 
Glauber's  salt  and  washing  soda  Na2COs,10H20,  have  a  vapor  ten- 
sion approaching  that  of  water  itself,  lose  their  water  at  ordinary 
temperatures  at  a  rapid  pace.  Now,  atmospheric  air  is  usually 
less  than  two-thirds  saturated  with  water  vapor,  and  the  partial 
pressure  (p.  72)  of  this  vapor  opposes  the  dissociation  and  tends  to 
prevent  the  liberation  of  the  water.  Thus  at  9°,  the  vapor  tension 
of  water  being  8.6  mm.,  the  average  vapor  pressure  of  water  in 
the  atmosphere  will  be  about  5  mm.  Any  hydrate  with  a  greater 
aqueous  tension  than  5  mm.  at  9°,  such  as  Glauber's  salt,  will 
therefore  decompose  spontaneously  in  an  open  vessel.  But  those 
with  a  lower  vapor  tension,  such  as  the  pentahydrate  of  cupric 
sulphate  with  a  tension  of  2  mm.  at  9°,  will  not  do  so.  Granular 
calcium  chloride  CaCl2,2H20  is  used  in  drying  gases  because  it 
has  an  exceedingly  low  tension  of  water  vapor,  and  combines  with 
water  vapor  to  form  CaCl2,6H20. 

The  water  of  hydration  is  known  colloquially  in  chemistry  as 
water  of  crystallization.  The  term  was  introduced  when  it  was  first 


WATER  97 

observed  that  a  hydrate,  in  decomposing,  crumbles  and  loses  its 
original  crystalline  form.  But  the  phrase  is  misleading.  Sulphur, 
potassium  chlorate,  and  thousands  of  other  substances  are  crys- 
talline, yet  do  hot  contain  the  elements  of  water.  All  pure  chemi- 
cal substances,  in  solid  form,  when  in  stable  physical  condition,  are 
crystalline.  Amorphous  (i.e.,  non-crystalline)  substances,  like  wax 
and  glass,  are  supercooled  liquids. 

How  Formulse  and  Equations  are  Obtained.  —  In  the  last 
few  pages  several  formulae  (e.g.,  of  hydrates)  and  several  new  equa- 
tions have  been  given.  How  do  we  know  what  to  set  down  in 
making  an  equation?  We  cannot  learn  this  by  simply  writing 
formulae  on  a  piece  of  paper.  In  each  case,  experiments  must  be 
made  in  the  laboratory.  For  example,  how  do  we  know  that  the 
common  hydrate  of  cupric  sulphate  has  the  formula  CuS04,5H20, 
and  not  CuS04,H20?  We  must  make  a  quantitative  experiment. 
We  weigh  a  porcelain  dish  or  crucible,  first  empty,  and  then  with 
a  little  of  the  hydrate.  Suppose  the  difference  in  weight  to  be  2.05 
g.  (=  weight  of  hydrate).  We  then  heat  the  dish  and  contents, 
until  the  water  is  driven  out,  and  weigh  again.  The  difference 
is  now  only  1.31  g.  (wt.  of  anhydrous  cupric  sulphate).  The 
water,  therefore,  weighed  2.05  -  1.31  =  0.74  g.  Assuming  that 
we  know  the  formulae  (compositions)  of  cupric  sulphate  and  of 
water,  we  obtain  their  formula-weights:  CuS04  =  63.57  +  32.06 
+  4  X  16  =  159.63;  andH20  =  2  X  1.008  +  16  =  18.016.  The 
formula  must  be  CuSO4,£H20.  Also 

159.63  :  x  X  18.016  ::  1.31  :  0.74. 

Solving  for  x,  we  have  x  X  18.016  X  1.31  =  159.63  X  0.74,  or 
x  =  159.63  X  0.74/18.016  X  1.31  =  5.00.  The  formula  is  there- 
fore CuS04,5H20,  and  the  equation  for  the  decomposition: 

CuS04,5H20  -*  CuS04  +  5H20. 

To  make  an  equation,  we  must  note  what  substances  are  taken, 
and  recognize  by  their  properties  all  the  substances  produced.  If 
all  the  substances  are  well  known,  and  we  can  find  their  formulae 
in  a  book,  we  can  make  the  equation  at  once.  If  we  cannot  find 
the  formulae,  we  make  measurements  to  determine  the  proportions 
by  weight,  calculate  the  formulas,  and  then  make  the  equation. 


98 


COLLEGE    CHEMISTRY 


Composition  of  Water.  —  The  proportion  of  hydrogen  to  oxy- 
gen, in  water,  by  weight,  is  2  :  15.879,  or  2.016  :  16.  The  propor- 
tion by  volume  is  2.0027  volumes  of  hydrogen  to  1  volume  of 
oxygen.  That  the  proportion  by  volume  is  very  close  to  2  :  1  may 
easily  be  shown  by  mixing  hydrogen  and  oxygen  in  this  propor- 
tion, in  a  strong  tube,  and  exploding  the  mixture  by  means  of  a 
spark  from  an  induction  coil.  The  resulting  steam  condenses  and 
the  whole  gas  vanishes.  If  different  proportions  are  used,  the 
excess  of  one  of  the  gases  remains  uncombined. 

Gay-Lussac's  Law  of  Combining  Volumes. 

—  The  almost  mathematical  exactness  with  which 
small  integers  express  this  proportion  is  not  a  mere 
coincidence.  Whenever  gases  unite,  or  gaseous  prod- 
ucts are  formed,  the  proportions  by  volume  (meas- 
ured at  the  same  temperature  and  pressure)  of  all 
the  gaseous  bodies  concerned  can  be  represented 
very  accurately  by  ratios  of  small  integers.  This  is 
called  Gay-Lussac's  law  of  combining  volumes  (1808). 
Thus,  when  the  above  experiment  is  carried  out 
at  100°,  in  order  that  the  product,  water,  may  be 
gaseous  also,  it  is  found  that  the  three  volumes  of 
the  constituents  give  almost  exactly  two  volumes 
of  steam.  For  example,  15  c.c.  of  hydrogen  and 
7.5  c.c.  of  oxygen  give  15  c.c.  of  steam.  Of  course 
the  hydrogen,  oxygen,  and  steam  must  be  measured 
at  the  same  pressure,  and  the  temperature  must 
remain  constant  (100°)  during  the  experiment. 
Proper  manipulation  secures  the  former,  and  a 
jacket  filled  with  steam  (Fig.  53)  the  latter  con- 
dition. Strips  of  paper,  1,  2,  and  3,  are  pasted  on 
the  jacket  in  such  a  way  that  equal  lengths  of 
the  eudiometer,  in  this  case  a  straight  one,  are 
laid  off.  The  three  divisions  being  filled  with  a  mixture  of  hydro- 
gen and  oxygen  in  the  proper  proportions,  the  gas,  after  the 
explosion,  shrinks  so  as  to  occupy,  at  the  same  pressure,  only  two 
of  them. 


Fia.  53. 


WATER  99 

Exercises.  —  1.   Name  some  familiar  transitions  (p.  86)  from 
one  physical  state  to  another. 

2.  What  evidence  is  there  in  the  common  behavior  of  ether 
and  chloroform  to  show  that  these  liquids  have  high  vapor  tensions? 

3.  If  the  pressure  of  the  steam  hi  a  boiler  is  ten  atmospheres, 
at  what  temperature  is  the  water  boiling  (p.  88)? 

4.  How  many  grams  of  water  could  be  heated  from  20°  to  100° 
by  the  heat  required  to  melt  1  kg.  of  ice  at  0°? 

5.  What  do  you  infer  from  the  fact  that  alum  and  washing  soda 
lose  their  water  of  hydration  when  left  in  open  vessels,  while 
gypsum  does  not? 

6.  Which  fact  shows  most  conclusively  that  hydrates  are  true 
chemical  compounds? 

7.  Gypsum  is  a  hydrate  of  calcium  sulphate  (CaS04).    If  6  g. 
of  gypsum,  when  heated,  lose  1.256  g.  of  water,  what  is  the  formula 
of  the  hydrate? 

8.  In  what  ways  does  a  hydrate  differ  from,  (a)  a  solution,  (6)  an 
hydroxide? 

9.  Should  you  expect  to  find  any  difference,  in  respect  to  chemi- 
cal activity,  between  the  three  forms  of  water?    Have  we  had  any 
experimental  confirmation,  or  the  reverse,  of  this  conclusion  (p. 
51)? 

10.  Name  some  crystalline  substances  which  are  not  used,  or 
do  not  occur  in  the  form  of  hydrates. 

11.  Define  the  purposes  for  which  evaporation  and  distillation 
are  used. 


CHAPTER  IX 
MOLECULAR  WEIGHTS  AND  ATOMIC  WEIGHTS 

GAY-LUSSAC'S  law  (p.  98)  shows  that,  when  substances  are 
measured  in  the  gaseous  condition,  and  by  volume  (not  by  weight), 
the  proportions  in  which  they  combine  can  be  represented  by  small 
whole  numbers,  such  as  2  :  1,  or  1  :  1,  or  2  :  3.  The  numbers  are 
much  simpler  than  when  proportions  by  weight  are  employed. 
Thus,  lead  and  oxygen  combined  in  the  proportions  100  :  7.72  by 
weight.  It  would  seem,  therefore,  that  the  shortest  route  to 
simple  methods  for  expressing  combining  proportions  must  lie 
through  a  study  of  volumes  of  gases  and  vapors. 

MOLECULAR  WEIGHTS 

The  Chemical  Unit  of  Volume  for  Gases:    22.4  Liters.— 

The  first  thing  we  require  is  a  suitable  unit  volume.  In  making  a 
choice,  we  have  to  keep  in  mind  the  fact  that  many  substances 
cannot  easily  be  converted  into  vapor,  and  that  therefore  meas- 
urement of  gaseous  volumes  cannot  entirely  displace  the  measure- 
ment of  weights.  The  measurement  of  gaseous  volumes  is  only 
to  furnish  the  key  to  the  system.  Hence,  in  choosing  our  unit 
volume  of  gas,  we  must  choose  one  which  bears  a  simple  relation 
to  our  units  of  weight.  Now,  the  unit  of  volume  chosen  is  that  of 
32  grams  of  oxygen,  which,  at  0°  and  760  mm.  pressure,  is  22.4  liters. 
At  this  stage,  it  may  appear  that  this  is  an  unduly  large  unit 

—  that  16  grams  of  oxygen,  occupying  11.2  liters,  might  have 
sufficed.    As  we  proceed,  however,  we  shall  find  that  a  smaller 
unit  than  22.4  liters  leads  to  a  number  less  than  unity  for  the 
atomic  weight  of  hydrogen.     There  is  no  theoretical  or  chemical 
objection  to  a  unit  involving  an  atomic  weight  for  hydrogen  that 
is  less  than  1,  but  chemists  are  unanimous  in  preferring  to  have  an 
arithmetical  unit  in  the  scale,  simply  as  a  matter  of  convenience. 
So,  reasoning  back  from  this  decision,  they  have  found  it  necessary 

—  we  shall  perceive  the  reason  presently  —  to  choose  22.4  liters  in 
the  gaseous  condition  as  the  unit  quantity  of  a  substance. 

100 


MOLECULAR  WEIGHTS 


101 


We  shall  understand  what  follows  much  more  readily  if  we  have 
before  us,  in  our  minds  at  least  or,  better  still,  in  the  form  of  a 
wooden  box,  a  representation  of  this  unit  volume  (Fig.  54).  A 
cube  11.1  inches  in  height  holds  22.4  liters/  :It*i«  to  be.,  under- 
stood that  under  conditions  other  than ... 
0°  and  760  mm.,  this  unit  volume, 
changes  in  accordance  with  the  laws 
of  gases.  In  this  way,  it  always  con- 
tains the  same  quantity  of  a  given  kind 
of  gas.  In  what  follows,  the  standard 
conditions  are  assumed,  unless  other 
conditions  are  specifically  mentioned. 


Q.M.V. 

22.4  LITEBS 


FIG.  54. 

Occupying    the    Unit    Volume  —  22.4   Liters.  — 

In  order  that  we  may  keep  in  touch  with  the  weights,  the  following 
table  gives  the  weights  of  equal  volumes  of  several  gases  and 
vapors.  The  first  column  contains  the  weights  of  1  liter.  The 
one-thousandth  part  of  each  of  these  weights'  is  the  density  * 
(p.  73)  of  the  gas  (weight  of  1  c.c.).  The  experimental  method 
of  measuring  the  weight  of  1  liter  of  a  gas  has  already  been  de- 
scribed (p.  73).  In  the  second  column  are  the  weights  of  22.4 
liters,  obtained  by  multiplying  the  values  in  the  first  column  by 
22.4.  It  will  be  observed  that  the  weights  of  equal  volumes  of 
the  gases  cover  a  wide  range  of  values  from  2  for  hydrogen  (col.  3) 
to  271.5  for  mercuric  chloride. 


Gases  or  Vapors. 

Weight*  of  One 
Liter,  0°  and  760 
mm. 

Weight  of  22.4 
Liters  (Molecular 
Weight). 

Hydrogen          

0.090 

2.016 

Oxvffen 

1.429 

32.00 

Chlorine                                           

3.166 

70.92 

Hydrogen  chloride                    ...... 

1.628 

36.468 

Carbon  dioxide                      

1.965 

44.00 

Water                                      

0.8045 

18.016 

Mercury                           

8.932 

200.6 

Mercuric  chloride       

12.097 

271.52 

Air                                                   

1.293 

28.955 

-   • 

*  Sometimes  density  is  expressed  on  the  basis  air  =  1.  One  liter  of  air 
weighs  1.293  g.  Hence,  if  one  liter  of  a  gas  weighs  3.6  g.,  its  density,  air  =  1, 
is  found  by  the  proportion:  1.293  :  1  ::  3.6  :  as. 


102  COLLEGE    CHEMISTRY 

The  values  for  the  vapors  of  water  (b.-p.  100°),  mercury  (b.-p. 
357°) ,  and  mercuric  chloride  (b.-p.  300°)  are  measured  at  high 
temperatures  and  reduced  by  rule  (pp.  71-72)  to  0°  and  760  mm. 

Molecular  Weights.  — In.  this  discussion  of  volumes  and  of 
weights,  we  must  not  o'verlgok  the  interpretation  of  our  results  in 
terms  of  molecules.  The  masses  of  gas  we  handle  are  aggregates 
of  molecules,  and  the  molecules  are  physically  the  real  units  of 
matter. 

Now,  according  to  Avogadro's  law,  equal  volumes  of  gases  (at 
the  same  t.  and  p.)  contain  equal  numbers  of  molecules.  The 
weights  in  each  column  of  the  table  are  therefore  weights  of  equal 
numbers  of  molecules.  The  chemical  units  in  the  last  column 
show,  therefore,  the  relative  weights  of  the  individual  molecules  of 
the  substances  named.  On  this  account  they  are  called  the  mo- 
lecular weights  of  the  respective  substances. 

Since  the  22.4-liter  volume  holds  32  grams  of  oxygen  and  2.016 
grams  of  hydrogen  —  the  gram  being  used  throughout  —  this 
volume  is  called  the  gram-molecular  volume  (G.M.V.)  and  the 
weights  just  mentioned  are  the  gram-molecular  weights.  Fre- 
quently, these  ponderous  terms  are  shortened  to  molar  volume  and 
molar  weight,  and  the  latter  even  to  mole.  Thus,  a  mole  of  chlo- 
rine is  70.92  g.  of  the  simple  substance  and  a  mole  of  hydrogen 
chloride  is  36.468  g.  of  the  compound.* 

Measurement  of  Molar  Weights  (Moles').  —  We  may  now 
state,  in  brief,  the  method  of  finding  the  molar  (gram-molecular) 
weight  of  a  substance  thus:  Weigh  a  known  volume  of  the  substance, 

*  A  common  question  is:  Do  not  molecules  of  different  substances  differ 
in  size,  and  will  not  the  numbers  required  to  fill  the  G.M.V.  therefore  be  differ- 
ent? The  answer  is  that  the  molecules  are  all  so  small  compared  with  the 
spaces  between  them  (at  760  mm.)  that  the  distances  from  surface  to  surface 
are  practically  the  same  as  from  center  to  center.  A  G.M.V.  of  oxygen, 
when  liquefied,  gives  less  than  32  c.c.  of  liquid  oxygen,  or  less  than  1/700  of 
the  volume  as  gas.  It  is  only  when  gases  are  so  severely  compressed  that  the 
nearness  of  the  molecules  to  one  another  approaches  that  found  in  the  liquid 
condition  that  the  effects  of  the  bulk  of  the  molecules  become  conspicuous, 
and  a  difference  in  the  behavior  of  different  gases  is  noticeable.  But  in  the 
work  discussed  in  this  chapter,  pressures  over  one  atmosphere  are  intention- 
ally avoided. 


ATOMIC   WEIGHTS  103 

at  any  temperature  and  pressure  at  which  it  is  gaseous,  reduce  this 
volume  by  rule  to  0°  and  760  mm.,  and  calculate  by  proportion  the 
weight  of  22.4  liters  (see  Exercises  1,  2,  3,  5). 

That  quantity  of  each  gaseous  substance  which  at  0°  and  760  mm. 
would  fill  the  G.M.V.  cube  is  the  unit  quantity  of  the  substance  for 
all  theoretical  purposes  in  chemistry.  It  represents  the  relative 
weight  of  the  molecules  of  the  substance.  We  shall  employ  it 
presently  for  the  purpose  of  determining  the  relative  weights  of 
atoms,  or  atomic  weights. 

The  Number  of  Molecules  in  a  Mole.  —  The  molecular 
weight  or  mole  of  a  substance  is  not  the  weight  of  a  single  molecule. 
It  is  only  the  relative  weight  of  the  molecule  of  the  substance.  It 
is,  however,  the  weight  in  grams  of  a  fixed  number  of  molecules, 
for  22.4  liters  (or  any  other  volume)  contains  equal  numbers  of 
molecules  of  different  gases.  The  actual  number  has  been  deter- 
mined. Thus  Jean  Perrin  found  values  by  several  experimental 
methods  which  ranged  between  5.9  X  1023  (that  is,  59  followed  by 
22  ciphers)  and  6.9  X  1023.  Rutherford,  using  an  entirely  differ- 
ent plan,  obtained  5.7  X  1023  for  the  gas  helium.  The  value 
which  is  accepted  as  most  accurate  was  that  obtained  by  R.  A. 
Millikan  of  the  University  of  Chicago,  by  the  use  of  a  still  differ- 
ent method,  namely,  6.07  X  1023  (or  60702i). 

ATOMIC  WEIGHTS 

Chemical   Unit   Quantities  of  Elements.  —  We  are  now 

approaching  the  question  of  units,  in  which  to  express  combining 
proportions,  from  a  different  view  point  from  that  employed  in  the 
earlier  chapters.  We  were  then  assigning  numbers  for  the  quan- 
tities of  the  constituent  elements  of  a  compound  (such  as  iron  : 
oxygen  ::  111.68  :  48,  p.  9)  without  any  consideration  of  the 
magnitude  of  the  total  weight  of  the  constituents.  At  that  time, 
we  had  no  reason  before  us  to  indicate  that  this  total  might  require 
consideration.  We  now  start  by  determining  and  assigning  the 
total  weight  of  the  compound,  and  it  is  our  next  task  to  consider 
the  subdivision  of  this  total  amongst  the  constituents.  Evidently, 
if  the  unit  quantity  of  the  compound  has  been  properly  chosen, 
it  must  be  subdivisible  into  one  or  more  unit  quantities,  of  suit- 
able dimensions,  of  each  element  in  the  compound.  Let  us  now 


104  COLLEGE    CHEMISTRY 

set  down,  and  examine  the  results  of  such  a  subdivision  in  the 
case  of  several  compounds.  To  be  more  precise,  we  take  22.4 
liters  of  every  substance  —  one  cubeful  in  the  gaseous  condition 
—  as  the  total  quantity.  We  make  an  analysis  of  a  sample  of 
the  substance,  in  case  it  is  a  compound,  to  ascertain  the  propor- 
tions in  which  the  elements  are  present  in  it.  We  then  divide 
the  weight  of  22.4  liters  of  the  compound  between  the  different 
elements  in  the  proportion  shown  by  the  analysis. 

Far  example,  the  cube  holds  36.468  g.  of  hydrogen  chloride  gas. 
This  amount,  when  decomposed,  yields  1.008  g.*  of  hydrogen  and 
35.46  g.  of  chlorine. 

Another  example:  Suppose  the  substance  is  a  liquid,  like  phos- 
phorus oxychloride.  We  determine  the  weight  of  a  measured 
volume  of  its  vapor,  at  a  properly  chosen  temperature  and  pres- 
sure, and  the  result  gives  us,  by  calculation,  the  weight  of  22.4 
liters,  the  molecular  weight,  viz.,  153.38.  That  is,  153.38  g.  of 
the  substance  would  fill  the  cube,  if  it  could  be  kept  as  vapor  at 
0°  and  760  mm.  The  analysis  shows  that  this  amount  of  the 
substance  contains  31  g.  of  the  element  phosphorus,  16  g.  of  the 
element  oxygen,  and  106.38  g.  of  the  element  chlorine. 

In  the  following  table  a  few  sample  results  of  the  process  just 
outlined  are  given.  The  first  column  contains  the  molar  weight, 
i.e.,  the  weight  of  the  substance  which  occupies  the  G.M.V.  cube. 
In  the  other  columns  are  entered  the  weights  of  the  various  ele- 
ments which  together  make  up  the  total  molar  weight.  To  sim- 
plify the  numbers,  the  values  used  are  hydrogen  1,  phosphorus 
31,  mercury  200,  instead  of  1.008,  31.04,  and  200.6,  respectively. 

Atomic  Weights.  —  To  contain  similar  data  for  all  the  volatile 
compounds  of  every  known  element,  a  huge  table,  of  which  this 
might  be  a  small  corner,  would  be  required.  With  such  a  table  at 
hand  the  atomic  weight  of  each  element  could  promptly  be  picked 
out.  Thus,  in  the  carbon  column  it  would  be  found  that  all  the 
weights  of  carbon  were  either  12  or  integral  multiples  of  12,  and 

*  It  will  be  observed  that  if  the  unit  for  molecular  weights  had  been  kss 
than  the  number  of  molecules  in  22.4  liters  of  oxygen,  then  an  equal  number  of 
molecules  of  hydrogen  chloride  would  have  contained  less  than  1.008  g.  of 
hydrogen,  and  the  atomic  weight  of  this  element  would  then  have  been  lees 
than  unity. 


ATOMIC   WEIGHTS 


105 


Substance. 

Molar 
Weight. 

Weights  of  Constituents  in  Molar  Weight. 

Hydrogen. 

| 

! 

! 

& 

Carbon. 

1 

Molecular 
Formula. 

Hydrogen  chloride.    .    .    . 
Chlorine  dioxide  
Phosphorus  trichloride.    . 
Phosphorus  oxychloride   . 
Phosphorus  pentoxide  .    . 
Phosphine  

36.46 
67.46 
137.38 
153.38 
284 
34 
18 
16 
26 
28 
30 
60 
235.46 
270.92 

1 

35.46 
35.46 
106.38 
106.38 

HCI 
C102 
PC13 
POC1, 

P4Ol0 

PH, 
H20 
CH4 
C2H2 
C2H4 
CH20 
C2H402 
HgCl 
HgCl2 

32 

'ie' 

160 

si 

31 
124 
31 

.... 

.... 

3 

2 
4 
2 
4 
2 
4 

Water  

.... 

16 

Methane    

12 
24 
24 
12 
24 

266' 
200 

Acetylene  

Ethylene   

Formaldehyde 

35^46 

16 
32 

Acetic  acid 

Mercurous  chloride    .    .    . 
Mercuric  chloride  .... 

70.92 

this  is  therefore  the  most  convenient  unit  weight  (and  therefore 
the  atomic  weight)  of  carbon.  Similarly,  the  atomic  weight  of 
oxygen  is  16,*  of  phosphorus  31,  of  mercury  200  (see  Exercise  4). 

The  fact  that  all  the  numbers  in  any  one  column  turn  out  to  be 
even  multiples  of  a  single  number  need  not  seem  mysterious.  The 
molecule  of  every  compound  containing  chlorine  must  contain 
one,  two,  three,  or  some  other  whole  number  of  chlorine  atoms,  for 
chlorine  atoms,  like  other  atoms,  do  not  furnish  fractions  of  atoms 
in  any  cases  of  combination.  Now,  the  weight  of  chlorine  in 
60702i  atoms,  assuming  one  atom  of  chlorine  to  each  molecule  in 
22.4  liters  of  some  gas  containing  chlorine,  must  be  35.46  g. 
Hence,  if  the  weight  of  chlorine  in  22.4  liters  (60702i  molecules) 
of  the  compound  differs  from  35.46  g.,  it  can  do  so  only  because 
there  are  two  atoms  of  chlorine  per  molecule,  giving  2  X  35.46  g., 
or  three  atoms  giving  3  X  35.46  g.  of  chlorine,  and  so  forth.  Thus 
the  quantities  of  chlorine  in  the  G.M.V.  of  all  compounds  of 
chlorine  must  be  a  multiple  of  35.46  by  unity  or  some  other  integer. 

When  the  atomic  weights  have  finally  been  selected,  we  can  go 
through  the  table  and  change  all  the  numbers  into  multiples  of  the 

*  The  difference  between  the  unit  quantity  of  oxygen  in  compounds 
(namely  16)  and  the  unit  quantity  of  free  oxygen  (32)  will  be  discussed  presently. 


106  COLLEGE    CHEMISTRY 

chosen  atomic  weights.  Thus,  for  70.92  we  write  2  X  35.46,  and 
for  106.38  we  write  3  X  35.46,  and  so  forth.  The  reader  should 
prepare  such  a  modification  of  the  table.  With  this  new  form  of 
the  table  before  us,  we  can,  finally,  replace  the  atomic  weights  by 
the  symbols  which  stand  for  them,  writing,  for  35.46,  Cl,  for 
2  X  35.46,  C12,  and  so  forth.  The  results  of  doing  this  in  each  line, 
i.e.,  for  each  substance,  are  collected  at  the  ends  of  the  lines  in  the 
last  column  of  the  table.  The  reader  should  himself  repeat  the 
substitutions  of  the  symbols,  and  so  verify  the  formulae  given. 
These  formulae,  since  they  are  based  on  the  molecular  weights, 
in  such  a  way  that  when  the  numerical  values  are  substituted  for 
the  symbols  the  total  restores  to  us  the  molecular  weight,  are 
called  molecular  formulae. 

As  a  definition,  the  atomic  weight  of  an  element  may  be  stated 
to  be:  The  smallest  of  the  weights  of  the  element  found  in  the 
molecular  weights  of  all  its  volatile  compounds,  so  far  as  these  have 
been  examined. 

It  is  hardly  necessary  to  add  that  the  atomic  weights,  found  as 
described  above,  are  equally  serviceable  in  expressing  the  compo- 
sitions of  compounds  which  are  not  volatile.  The  atoms  in  non- 
volatile compounds  are  identical  in  properties  with  the  atoms  of 
the  same  elements  in  volatile  compounds.  If  an  element  gives  no 
volatile  compounds,  other  methods  of  fixing  its  atomic  weight  are 
available  (see  Dulong  and  Petit's  law,  p.  108). 

Although  in  this  section,  as  well  as  elsewhere,  we  have  empha- 
sized the  fact  that  atoms  are  not  divided  into  parts,  this  must  not 
be  taken  to  mean  that  atoms  are  incapable  of  being  broken  up. 
It  means  only  that  in  ordinary  chemical  changes,  the  atoms  com- 
bine and  separate  as  wholes.  Indeed,  we  now  know  that  the  atom 
of  radium  (q.v.)  gives  off  atoms  of  helium,  and  leaves  an  atom  of 
lead,  and  that  the  atoms  of  one  or  two  other  elements  disintegrate 
in  a  similar  way.  Some  day  means  of  breaking  up  any  or  all 
kinds  of  atoms  may  be  discovered. 

Many  chemists  have  contributed  to  the  determination  and  re- 
vision of  the  atomic  weights.  The  Swedish  chemist,  Berzelius, 
devoted  many  years  to  the  accurate  measurement  of  combining 
proportions.  Stas,  a  Belgian  (1860-1870),  made  a  number  of 
determinations  with  great  exactness.  Morley's  (1895)  value 
for  combining  proportions  of  hydrogen  and  oxygen  alone  repre- 


ATOMIC   WEIGHTS  107 

sented  several  years  of  work.  T.  W.  Richards  of  Harvard  Uni- 
versity has  recently  carried  many  of  the  values  to  a  higher  degree 
of  accuracy. 

Why  22.4  Liters  was  Chosen  as  the  Unit  Volume.  —  We 

can  now  see  why  the  volume  occupied  by  32  g.  of  oxygen,  namely, 
22.4  liters,  was  taken  as  the  standard  for  the  scale  of  molecular 
quantities.  This  gave  us,  for  example,  36.468  g.  as  the  weight  of 
22.4  1.  of  hydrogen  chloride,  which  in  turn  contains  1.008  g.  of 
hydrogen.  A  smaller  weight  of  oxygen,  with  correspondingly 
smaller  standard  volume,  would  have  held  an  amount  of  hydrogen 
chloride  (and  of  other  compounds  containing  one  atom  of  hydro- 
gen per  molecule)  which  would  have  been  less  than  1  gram.  The 
choice  was  made  to  secure  something  close  to  an  arithmetical 
unit  in  the  scale. 

Advantages  of  Atomic  Weights.  —  Although  the  method 
of  selecting  atomic  weights  involves  rather  complex  reasoning, 
these  weights  repay  the  trouble,  because  they  represent  the  rela- 
tive weights  of  the  atoms  themselves.  They  are  thus  much  more 
valuable  in  helping  us  to  understand  chemical  behavior  and  in 
enabling  us  to  classify  the  phenomena  of  chemistry  than  would  be 
any  other  units  of  weight  we  might  have  chosen.  The  following 
are  some  of  the  advantages  they  offer: 

1.  The  atomic  weight  of  an  element  has  but  one  value,  and 
this  value  is  definitely  determinable.     The  advantages  of  using 
Avogadro's  principle  (1811),  and  taking  a  unit  volume  of  gas  as 
the  basis  of  chemical  units,  were  not  perceived  by  chemists  until 
Cannizzaro,  in  1858,  succeeded  in  setting  them  forth  in  a  con- 
vincing manner.     Previous  to  that  time,  different  chemists  used 
different  unit  weights  for  the  same  element,  and  therefore  assigned 
different  formulae  to  the  same  compound,  and  much  confusion  ex- 
isted.    After  1858  chemists  united  upon  the  present  values  for 
atomic  weights. 

2.  The  atomic  weight  of  an  element  has  a  valence  (p.  61),  while 
equivalents  are  equi-valent.     While  valence  is  a  helpful  conception 
in  all  branches  of  chemistry,  organic  chemistry  is  especially  in- 
debted to  the  conception  of  the  quadrivalence  of  carbon  for  much 
of  its  development  and  most  of  its  organization.     The  full  illus- 
tration of  this  point  is  beyond  the  limits  of  the  present  book. 


108 


COLLEGE    CHEMISTRY 


3.  The  periodic  system  (q.v.),  the  basis  of  a  plan  for  classifying 
the  properties  of  all  chemical  substances,  is  founded  upon  the 
atomic  weights. 

4.  Dulong  and  Petit's  law  is  based  upon  atomic  weights.     This 
law  furnishes  also  an  alternative  means  of  determining  atomic 
•weights  that  has  frequently  rendered  valuable  service,  and  on  this 
account  forms  the  subject  of  the  next  section. 

Dulong  and  Petit's  Law,  an  Alternative  Means  of  Deter- 
mining Atomic  Weights.  —  It  was  first  pointed  out  (1818)  by 
Dulong  and  Petit,  of  the  Ecole  Polytechnique  in  Paris,  that  when 
the  atomic  weights  of  the  elements  were  multiplied  by  the  specific 
heats  of  the  simple  substances  in  the  solid  condition,  the  products 
were  approximately  the  same  in  all  cases.  In  other  words,  the  spe- 
cific heats  are  inversely  proportional  to  the  magnitudes  of  the 
atomic  weights.  The  table,  in  which  round  numbers  have  been 
used  for  the  atomic  weights,  shows  that  the  product  lies  usually 
between  6  and  7,  averaging  about  6.4: 


Element. 

Atomic 
Wt. 

Sp.  Ht. 

Prod- 
uct. 

Element. 

Atomic 
Wt. 

Sp.  Ht. 

Prod- 
uct. 

Lithium  .    .    . 

7 

0.94 

6.6 

Iron     

56 

0.112 

6.3 

Sodium  .    .    . 

23 

0.29 

6.7 

Zinc    

65.4 

0.093 

6.1 

Magnesium  . 

24.3 

0.245 

6.0 

Bromine  (Solid) 

80 

0.084 

6.7 

Silicon    .    .    . 

28.3 

0.16 

4.5 

Gold    

197 

0.032 

6.3 

Phosphorus 

Mercury  (Solid) 

200 

0.0335 

6.7 

(Yellow)    . 

31 

0.19 

5.9 

Uranium    .    .    . 

238.5 

0.0276 

6.6 

Calcium.    .    . 

40 

0.170 

6.8 

Another  way  of  expressing  this  law  will  give  it  greater  chemical 
significance.  The  specific  heats  are  the  amounts  of  heat  required 
to  raise  one  gram,  that  is  one  physical  unit,  of  each  element  through 
one  degree.  When  we  multiply  this  by  the  atomic  weight,  we 
obtain  the  amount  of  heat  required  to  raise  one  gram-atomic  weight 
of  the  element,  that  is,  one  chemical  unit,  through  one  degree. 
The  values  of  this  product  are  approximately  equal.  Since  there 
are  equal  numbers  of  atoms  in  one  gram-atomic  weight  of  each  ele- 
ment, it  follows  that :  Equal  amounts  of  heat  raise  equal  numbers 
of  atoms  of  all  elements  in  the  solid  form  through  equal  intervals  of 
temperature. 


MOLECULAR  EQUATIONS  109 

It  will  be  seen  at  once  that,  although  the  law  of  Dulong  and 
Petit  is  purely  empirical,  it  may  nevertheless  be  used  for  fixing  the 
atomic  weight  of  an  element  of  which  no  volatile  compounds  are 
known.  We  can  always  measure  that  weight  of  such  an  element 
which  combines  with  one  atomic  weight  of  another  element.  Since 
the  elements  concerned  must  combine  atom  for  atom,  or  in  some 
simple  ratio  such  as  1  :  2  or  2  :  3,  it  follows  that  the  weight  found 
is  either  the  atomic  weight  or  some  multiple  or  submultiple  of  it 
by  a  whole  number.  When,  therefore,  we  multiply  this  weight 
by  the  specific  heat,  we  discover  at  once  whether  the  product  is 
6.4  or  some  simple  fraction  or  multiple  of  this  number.  For  ex- 
ample, suppose  the  atomic  weight  of  calcium  to  be  unknown. 
We  find  by  analysis  that  calcium  chloride  contains  20  parts  of 
calcium  combined  with  35.46  parts  (the  atomic  weight)  of  chlo- 
rine. Now  the  specific  heat  of  solid,  metallic  calcium  is  0.170. 
This  number  multiplied  by  20  gives  as  the  product  3.4.  Evi- 
dently, therefore,  the  atomic  weight  is  not  20,  but  40,  for  the 
product,  6.8,  then  agrees  fairly  well  with  the  average  for  other 
elements. 

MOLECULAR  FORMULAE 

• 

Molecular  Formulse  of  Compounds.  —  If  the  molar  formulae 
in  the  table  (p.  105)  be  examined  it  will  be  observed  that  several 
are  not  in  their  simplest  terms.  Thus,  the  formula  of  acetylene  is 
C2H2.  The  formula  CH  would  represent  the  composition  of  the 
substance  equally  well,  for  12  :  1  is  the  same  as  24  :  2.  But  the 
formula  CH  gives  a  total  of  only  13,  while  C2H2  shows  the  total 
weight  of  the  molecule  to  be  26  and  records  for  us  therefore  the 
weight  of  the  G.M.V.,  as  well  as  the  composition  of  the  substance. 
We  shall  find  this  additional  property,  peculiar  to  the  molecular 
formula,  to  be  a  feature  of  the  greatest  practical  value.  Some 
of  the  practical  uses  of  this  improvement  in  our  formulae  will  be 
illustrated  in  this  chapter,  and  there  is  an  example  of  one  of  them 
in  the  table  itself.  Thus,  the  molecular  formula  of  acetic  acid  is 
C2H4O2,  and  not  the -simpler,  identical  proportion  CH20.  The 
latter  is  the  molecular  formula  of  a  totally  different  substance, 
formaldehyde,  now  much  used  as  a  disinfectant.  The  vapor  of 
this  substance  has  only  half  the  density  of  acetic  acid  vapor,  and 
this  fact,  recorded  in  the  formula,  helps  to  remind  us  that  the 


110  COLLEGE    CHEMISTRY 

substances  are  different.  Still  another  substance  of  the  same 
composition  is  grape  sugar  (dextrose),  CeHisOe.  In  addition  to 
this  and  other  practical  advantages,  molecular  formulae  satisfy  also 
the  claim  of  logical  consistency.  If  the  symbols  Represent  the 
atomic  weights,  the  formulas  should  be  constructed  so  as  to  rep- 
resent the  molecular  weights. 

Molecular  formulae  like  C2H2  and  C2H402  are  easily  interpreted 
in  terms  of  the  atomic  hypothesis.  C  represents  one  atom  of 
carbon  and  H  one  atom  of  hydrogen.  But  there  is  no  reason 
why  a  molecule  of  acetylene  should  not  contain  two  atoms  of  each 
kind.  Similarly,  the  molecule  of  formaldehyde  contains  four  atoms 
(CH20),  and  one  of  acetic  acid  eight  atoms  (C2H4O2),  and  one  of 
dextrose  twenty-four  atoms  (Cel^Oe),  although  the  relative  num- 
bers of  each  kind  are  the  same.  Indeed  this  hypothesis  helps  to 
clear  the  matter  up,  for  chemists  go  so  far  as  to  account  for  the 
chemical  behavior  of  the  substances  by  an  imagined  geometrical 
arrangement  of  the  atoms  in  their  molecules,  and  these  three 
kinds  of  molecules  are  supposed  to  differ  hi  structure  as  well  as  in 
the  number  of  atoms  they  contain. 

The  Molecular  Weights  and  Formulae  of  Elementary 
Substances.  —  The  following  table  gives  the  densities  of  some 
elementary  substances,  including  those  of  which  the  substances 
previously  discussed  are  compounds.  The  first  column  shows  the 
atomic  weight,  which  in  each  case  is  the  minimum  weight  of  the 
element  found  in  a  G.M.V.  of  any  compound.  For  example,  16  g. 
of  oxygen  and  35.46  g.  of  chlorine  are  the  weights  in  the  amounts 
of  water  vapor  and  hydrogen  chloride,  respectively,  which  fill  the 
cube  (22.4  liters).  The  symbol,  in  the  next  column,  stands  for 
this  quantity  and  occurs  in  many  formulae,  such  as  H20  and  HC1. 
It  represents  the  combining  unit  or  atom.  In  the  third  column  is 
given  the  weight  of  the  free,  elementary  substance  which  fills  the 
G.M.V.  and  is  the  molecular  weight.  It  shows  the  weight  of  the 
molecule  relative  to  the  weights  of  the  other  molecules  in  the  same 
column,  and  to  the  weights  of  the  atoms  in  the  first  column.  In 
the  last  two  columns  are  given  the  molecular  weights  resolved 
into  multiples  of  the  atomic  weights  and  the  corresponding 
formulas. 


MOLECULAR  EQUATIONS 


111 


Element. 

Atomic 
Weight. 

Sym- 
bol. 

Weight  in 
G.M.V. 

Weight  in 
G.M.V.  Fac- 

torized. 

Formula 
of  Free 
Element. 

Oxygen          .    .    . 

16  00 

o 

32  00 

ovifi  no 

O« 

Hydrogen  

1  008 

H 

2  016 

2v  i    nno 

\J2 

H« 

Chlorine    

35  46 

Cl 

70  92 

2x35  46 

Phosphorus  

31  04 

p 

124  16 

4X31  04 

P- 

JMercury 

200  6 

Hg 

200  6 

i  y  900  fi 

Ho- 

Ozone 

16  00 

0 

48  00 

3x16  00 

"g 
O, 

Cadmium              .    . 

112  4 

Cd 

112  4 

1x112  4 

Cd 

Potassium  

39  10 

K 

39  10 

1X39  10 

K 

Sodium  

23  00 

Na 

23  00 

1X23  00 

Na 

Zinc    

65.37 

Zn 

65  37 

1X65  37 

Zn 

The  reader  cannot  fail  to  note  a  striking  peculiarity.  In  the  case 
of  chlorine  the  molecular  weight  is  70.92,  while  the  atomic  weight 
is  35.46.  With  hydrogen  and  oxygen,  also,  the  molecular  weight 
contains  two  atomic  weights.  Yet  this  is  not  a  general  rule,  for 
with  mercury  and  several  other  elements  the  molecular  and  atomic 
weights  are  alike,  while  with  phosphorus  the  molecular  is  four  times 
the  atomic  weight.  Evidently  there  is  no  rule,  and  each  element 
has  to  be  subjected  to  separate  experimental  study.  The  result  is 
that  for/ree,  elementary  chlorine  we  use  the  molecular  formula  C12,  for 
free  hydrogen  H2,  for  elementary,  uncombined  oxygen  the  formula  02. 
For  a  substance  like  phosphorus,  which  is  not  a  gas  and  is  not  often 
used  as  a  vapor,  the  formula  P  is  commonly  employed  by  chemists, 
to  avoid  the  larger  coefficients  which  ?4  introduces  into  equations, 
although  theoretically  the  latter  formula  would  be  the  strictly  cor- 
rect one. 

The  case  of  oxygen  demonstrates  clearly  the  necessity  of  using 
molecular  formulae,  even  for  simple  substances.  The  table  shows 
two  substances  containing  nothing  but  oxygen.  Ozone  (q.v.)  has 
a  molecular  weight  48,  being  a  gas  exactly  one-half  heavier  than 
ordinary  oxygen.  Its  formula,  therefore,  is  Os,  while  that  of  oxy- 
gen is  02.  Oxygen  and  ozone  are  entirely  different  chemical  indi- 
viduals. The  latter  has,  for  example,  a  strong  odor  and  is  much 
more  active.  Thus  polished  silver  remains  bright  indefinitely  in 
pure  oxygen,  but  oxidizes  quickly  when  placed  in  ozone. 

To  avoid  a  common  error,  the  reader  should  note  that  to  learn  the 
atomic  weight  of  an  element,  we  do  not  measure  the  molecular 


112  COLLEGE   CHEMISTRY 

weight  of  the  simple  substance.  The  molecular  weight  of  the  ele- 
mentary substance  may  be  a  multiple  of  the  atomic  weight,  and  we 
find  out  whether  it  is  such  a  multiple  only  after  the  atomic  weight 
has  been  determined.  The  atomic  weight  is  the  unit  weight  used 
in  compounds,  and  can  be  ascertained  only  by  a  study  of  com- 
pounds. The  molecular  weight  of  the  free  element  gives  us  only 
a  value  which  we  know  must  be  a  multiple  of  the  atomic  weight,  by 
1  or  some  other  integer.  Mol.  Wt.  =  At.  Wt.  X  x,  where  a;  is  1  or 
some  other  integer. 

Further  Discussion  of  the  Molecular  Formulae  of  Ele- 
mentary Substances.  —  Some  further  explanation  may  be  re- 
quired, to  the  end  that  the  reader  may  be  reconciled  to  accepting 
the  formulae  C12,  O2,  and  so  forth.  In  the  first  place,  he  should 
note  how  these  formulae  arose.  If  we  accept  Avogadro's  law,  and 
the  inference  from  it  to  the  effect  that  the  weights  of  equal  vol- 
umes of  gases  are  in  the  same  ratio  as  the  weights  of  their  indi- 
vidual molecules,  .then  we  cannot  escape  the  conclusion  to  which 
measuring  the  relative  densities  of  free  chlorine  and  hydrogen 
chloride,  for  example,  leads.  The  ratio  of  their  densities  is  70.92  : 
36.46.  That  is  to  say,  the  relative  weights  of  a  molecule  of  chlo- 
rine and  a  molecule  of  hydrogen  chloride  stand  in  this  ratio.  The 
molecule  of  chlorine  is  nearly  twice  as  heavy  as  the  molecule  of  the 
compound,  and  there  cannot  therefore  be  a  whole  molecule  of  chlorine 
in  a  molecule  of  hydrogen  chloride.  In  fact,  we  perceive  at  once 
that  the  molecule  of  hydrogen  chloride  must  contain  only  half  a 
molecule  of  chlorine  (35.46),  together  with  half  a  molecule  of 
hydrogen  (1).  In  other  words,  if  the  molecule  of  free  chlorine 
were  to  be  taken  as  the  atom  of  the  element,  then  the  molecule  of 
hydrogen  chloride  would  contain  only  half  an  atom  of  chlorine, 
which  would  be  contrary  to  our  definition  to  take  as  atoms  quan- 
tities which  are  not  divided.  So  we  choose  the  other  horn  of  the 
dilemma,  and  say  that  the  specimen  of  chlorine  in  the  molecule  of 
hydrogen  chloride  is  a  whole  atom  and  that  therefore  the  amount 
of  chlorine  in  the  molecule  of  free  chlorine  is  two  atoms,  and  its 
formula  C12.  Similarly,  the  weight  of  hydrogen  in  the  molecule  of 
hydrogen  chloride  is  1.008,  while  that  of  the  molecule  of  hydrogen 
is  2.016,  so  that  there  are  two  atoms  in  the  molecule  of  free  hydro- 
gen and  its  formula  is  H2.  Reasoning  in  like  manner  from  the 


MOLECULAR   EQUATIONS  113 

molecular  weights  of  oxygen  (32)  and  water  (18)  we  reach  the 
conclusion  that  the  molecule  of  oxygen  is  diatomic  (62). 

The  simple  fact  that  hydrogen  and  oxygen,  when  mixed,  do  not 
combine  (p.  59)  may  assist  in  reconciling  us  to  the  diatomic  nature 
of  their  molecules.  Some  part  of  the  mixture  has  to  be  heated 
strongly  to  start  the  interaction.  Now  the  molecular  formulae,  H2 
and  02,  suggest  that  each  gas  is  really  in  combination  already  (with 
itself),  and  they  therefore  explain  to  some  extent  the  indifference  of 
the  gases  towards  one  another.  If  the  molecules  were  free  atoms, 
they  could  not  encounter  one  another  continually  as  they  move 
about,  and  yet  escape  combination  as  we  observe  that  they  do. 
We  may  imagine  that  the  primary  effect  of  heating  is  to  decom- 
pose some  of  the  molecules,  and  liberate  hydrogen  and  oxygen  in 
the  atomic  condition,  and  that  the  combination  of  these  atoms 
starts  the  explosion  of  the  whole  mass. 

In  the  case  of  hydrogen,  the  diatomic  nature  of  the  molecules 
has  been  demonstrated  by  an  entirely  different  method  by  Lang- 
muir.  It  has  long  been  known  that  the  conductivity  of  hydrogen 
for  heat  is  greater  than  that  of  any  other  elementary  gas.  Thus, 
a  wire  raised  to  a  white  heat  in  air  by  means  of  an  electric  current, 
cannot  be  kept  at  a  red  heat,  even,  by  the  same  current  in  hydro- 
gen. In  other  gases,  heat  from  the  hot  wire  is  used  up  in  accel- 
erating the  motion  of  the  molecules  of  the  gas.  Langmuir  has 
shown,  however,  that  in  hydrogen,  additional  heat  is  consumed  in 
causing  decomposition  of  many  of  the  diatomic  molecules  into 
single  atoms: 


He  has  measured  the  percentage  of  molecules  dissociated  (at  760 
mm.),  and  found  that  it  varies  from  0.33  per  cent  at  2000°  to  13  per 
cent  at  3000°  and  34  per  cent  at  3500°.  When  the  temperature 
falls,  the  atoms  re-combine  to  diatomic  molecules.  It  may  also 
assist  in  making  the  matter  clear  if  we  note  that  the  atomic  weight 
of  an  element  is  the  unit  quantity  of  that  particular  variety  of 
matter,  when  it  is  in  combination.  The  unit  quantity  of  the  same 
variety  of  matter,  when  in  the  free  state,  as  a  substance,  need  not 
be  the  same.  We  should  not  expect  it  to  be  smaller,  but  it  might 
easily  be  twice  or  more  times  as  large. 


114 


COLLEGE    CHEMISTRY 


APPLICATIONS 

Applications:  Interactions  Between  Gases.  —  According  to 
Avogadro's  hypothesis,  if  we  filled  a  succession  of  vessels  of  equal 
dimensions  with  different  gases,  and  could  arrest  the  motion  of 
the  particles  and  observe  their  disposition,  we  should  find  that  the 
average  distance  from  particle  to  particle  would  be  the  same  in 
all  cases.  This  would  be  true  whether  our  vessels  were  filled  with 
single  gases,  with  homogeneous  mixtures,  or  with  gases  in  layers. 
Such  being  the  case,  if  any  chemical  change  is  brought  about  in  the 
mass  which  results  in  a  multiplication  of  the  molecules,  it  is  evi- 
dent that  the  volume  will  have  to  increase  in  order  that  the  spacing 
may  remain  the  same  as  before.  If  any  chemical  action  results  in 
a  diminution  of  the  number  of  molecules,  then  a  shrinkage  must 
take  place  in  order  that  the  spacing  may  be  preserved  as  before. 
Thus,  in  a  mixture  of  hydrogen  and  oxygen,  according  to  our 
hypothesis,  when  the  interaction  occurs,  the  following  change 
takes  place  between  neighboring  molecules: 

HH  +  00  +  HH  becomes  HOH  +  HOH. 

Since  the  oxygen  molecules,  which  form  a  third  of  the  whole,  dis- 
appear into  the  molecules  of  hydrogen,  the  tendency  to  preserve 
spacing  results  in  a  diminution  of  the  volume  by  one-third  (p.  98) . 
Thus  Gay-Lussac's  law  would  have  followed  as  a  natural  infer- 
ence from  Avogadro's  law,  if  the  former,  being  more  obvious,  had 
not  been  discovered  first. 

If  each  of  the  following  squares  represents  a  small  volume  con- 
taining 1000  molecules  of  gas,  then  2000  molecules  of  hydrogen 
and  1000  molecules  of  oxygen  give  2000  molecules  of  water  vapor. 
We  may  note,  in  passing,  that,  since  each  molecule  of  water  must 
contain  at  least  one  atom  of  oxygen,  at  least  2000  atoms  of  oxygen 
were  required,  and  must  have  been  furnished  by  the  1000  mole- 
cules of  oxygen.  Each  of  these  molecules  must  therefore  have 
split  into  two  atoms. 


This  method  of  looking  upon  chemical  interactions  between 
gases  gives  us  the  nearest  sight  which  we  can  have  of  the  behavior 


APPLICATIONS   OF   MOLECULAR  EQUATIONS  115 

of  the  molecules  themselves.  We  cannot  perceive  the  individual 
molecules,  but,  in  consequence  of  the  spatial  arrangement  which 
they  observe,  the  change  in  the  whole  volume  of  a  large  aggre- 
gate of  molecules  enables  us  to  draw  conclusions  at  once  in  regard 
to  the  behavior  of  the  single  molecules  in  detail. 

Applications:  Molecular  Equations.  —  To  utilize  the  fore- 
going considerations,  chemists  always  employ  in  their  equations  the 
molecular  formulae  for  the  gases  and  the  easily  vaporized  substances 
concerned.  Thus  far^  we  have  used  the  equation: 

2H      +     O    ->    H2O 
WEIGHTS:  2  X  1.008  16  18.016 

and  the  information  it  contained  was  exhausted  when  we  had  placed 
below  the  symbols  the  weights  for  which  they  stood.  But  the 
molecular  equation  is  much  more  instructive.  The  following 
shows  the  interpretations  to  which  the  molecular  equation  is  subject: 

2H2  +        02     ->      2H20 

WEIGHTS:         2  X  2.016  g.  32  g.  2  X  18.016  (=  36.032)  g. 

VOLUMES:          2  X  22.4  1.  22.4  1.          2  X  22.4  1. 

MOLECULES:  2  12 

The  weights,  although  doubled,  show  the  same  proportions,  so 
that  questions  of  weight  are  answered  as  easily  as  before.  These 
weights,  however,  being  molecular  weights,  or  multiples  thereof, 
can  be  translated  at  once  into  volumes,  and  questions  about  volumes 
can  also  be  answered.  Finally,  the  relative  numbers  of  each  kind  of 
molecules  can  be  read  from  this  equation,  for  the  coefficients  in 
front  of  the  formulae  represent  these  numbers.  Where  no  coef- 
ficient is  written,  1  is  to  be  understood.* 

Applications:     The  Making  of  Molecular  Equations.— 

To  make  a  molecular  equation,  we  first  make  an  equation  accord- 
ing to  the  rules  already  explained  (p.  51).  An  equation  like  that 
given  for  the  interaction  of  potassium  on  water  (p.  50)^«iC  +  H2O 
— >  KOH  +  H,  is  the  result.  Then  we  adjust  the  equation  so 

*  The  application  of  these  properties  of  molecular  equations  is  illustrated 
in  Chap.  XI  (pp.  149-153).  If  desired,  these  applications  may  be  taken  up 
after  the  next  section. 


116  COLLEGE    CHEMISTRY 

that  molecular  formulae  are  used  throughout.  The  hydrogen  must 
appear  as  H2,  or  a  multiple  of  this,  in  such  equations.  Hence 
the  whole  equation  must  be  multiplied  by  2: 


2K  +  2H20  ->  2KOH  -}>  H2 

Again,  the~equation  for  the  preparation  of  oxygen  from  potassium 
chlorate:  KC103  ->  KC1  +  30  (p.  27),  becomes: 
2KC1O3  ->  2KC1  +  3O2. 

Every  equation  containing  an  odd  number  of  atoms  of  a  substance 
whose  molecules  are  diatomic  must  be  multiplied  by  2.  Again, 
mercuric  oxide  decomposes  to  give  mercury  vapor  and  oxygen 
(p.  14),  and  the  molecules  of  mercury  are  monatomic  and  those  of 
oxygen  diatomic,  so  we  write: 

2HgO  -» 2Hg  +  O2. 

Finally,  the  formulae  of  substances  which  are  solid  or  liquid,  and 
cannot  be  easily  vaporized,  are  written  in  the  simplest  terms. 
Thus,  since  substances  like  the  copper  in  the  following  equation 
are  involatile,  the  molecular  weights  of  such  substances  are  un- 
known, and  their  molecular  formulae  likewise:  2Cu  -f  02  — »  2CuO. 
Furthermore,  in  the  case  of  substances  which  can  be  volatilized, 
although  the  molecular  weights  and  molecular  formulae  may  there- 
fore be  known,  we  do  not  usually  employ  the  molecular  formulae  if 
the  substance  is  not  used  in  the  form  of  vapor  in  the  laboratory. 
Thus,  the  molecular  formula  of  phosphorus  pentoxide  is  P4Oi0 
(p.  105).  But  we  generally  make,  and  use,  only  the  solid  form, 
and  not  the  vapor,  in  actual  work.  Hence  the  action  with  water 
is  usually  written  as  we  have  given  it  (p.  94),  rather  than  in  the 
form:  P4Oio  +  6H2O  ->  4H3PO4. 

Molecular  equations  will  be  used  exclusively  hereafter. 

Applications:  To  Cases  of  Dissociation.  —  Several  gases 
or  vapors  yield  smaller  values  for  their  densities,  and  therefore 
molecular  weights,  when  the  densities  are  measured  at  higher  tem- 
peratures. This  indicates  that  the  molecules  have  become  lighter, 
and  can  only  mean  that  decomposition  has  taken  place  in  conse- 
quence of  the  heating.  Behavior  of  this  kind  is  shown  both  by 
compounds  and  by  simple  substances. 


APPLICATIONS   OF   MOLECULAR  EQUATIONS  117 

For  example,  phosphorus  pentachloride  PC15,  although  a  solid, 
can  be  converted  into  vapor  without  much  difficulty.  Its  molec- 
ular weight,  if  it  underwent  no  chemical  change  during  the  vola- 
tilization, would  be  31  +  177.3  =  208.3.  The  density  actually 
observed  at  300°  and  760  mm.  pressure  gives  by  calculation  not 
much  more  than  half  this  value.  The  direct  inference  from  this 
is  that  the  molecules  have  only  half  the  (average)  weight  that  we 
expected;  or,  in  other  words,  are  twice  as  numerous  as  we  expected. 
The  explanation  is  found  when  we  examine  the  nature  of  the  vapor 
more  closely.  We  find  that  it  is  a  mixture  of  phosphorus  trichlo- 
ride and  free  chlorine,  resulting  from  a  chemical  change  according 
to  the  equation:  PC15  <=*  PC13  +  C12.  The  low  value  of  the  den- 
sity thus  tells  us  that  dissociation  has  taken  place.  From  the 
value  of  the  density  at  various  temperatures,  we  may  even  calcu- 
late the  proportion  of  the  whole  material  which  is  dissociated.  At 
300°  it  is  97  per  cent;  at  250°,  80  per  cent;  and  at  200°,  48.5  per 
cent.  Thus,  when  the  temperature  is  lowered,  progressive  re- 
combination takes  place  and  the  proportion  dissociated  becomes 
less.  Finally  the  vapor  condenses  and  yields  the  original  solid. 

Again,  sulphur  boils  at  445°,  but  can  be  vaporized  at  a  tempera- 
ture as  low  as  193°,  under  very  low  pressure.  At  this  temperature 
the  density  of  the  vapor  gives  the  molecular  weight  256  (=  8  X 
32),  and  the  molecular  formula  Sg.  That  is  to  say,  the  G.M.V. 
holds  256  g.  of  the  vapor  at  193°.  At  800°,  however,  the  density 
is  only  one-fourth  as  great,  and  the  G.M.V.  holds  only  64  g.  (82). 
This  means  that  256  g.  now  occupy  four  times  as  large  a  volume 
as  before,  and  the  increase  is  additional  to  the  effect  of  the  mere  ther- 
mal expansion,  which  is  allowed  for  in  the  calculation  and  elimi- 
nated. Hence  the  molecules  have  dissociated.  At  1700°  the 
molecular  formula  is  still  S2,  so  that  this  represents  the  limit  of 
dissociation:  Sg  «=*  4S2.  When  the  vapor  is  cooled,  the  density 
increases  once  more  and  at  193°  recovers  completely  the  greater 
value.  Similar  observations  show  that  phosphorus  vapor  at  313° 
is  all  P4,  but  at  1700°  one-half  of  the  molecules  are  P2.  Iodine 
vapor,  up  to  700°,  is  all  I2.  Beyond  this  temperature  the  density 
diminishes,  and  when  1700°  is  reached  the  vapor  is  all  I.  Thus 
the  molecules  are  diatomic  at  low  temperatures  and  monatomic  at 
high  ones.  The  densities  of  oxygen,  hydrogen,  and  chlorine  are 
not  measurably  affected  by  heating  to  1700°,  so  that  their  dia- 


. 

118  COLLEGE    CHEMISTRY 

tomic  molecules  exist  from  temperatures  far  below  0°  up  to 
1700°,  and  are  evidently  very  stable.  For  observations  on  hydro- 
gen above  1700°,  however,  see  p.  113. 

Applications:  Finding  the  Atomic  Weight  of  a  New  Ele- 
ment.—  By  way  of  reviewing  the  principles  explained  in  this 
chapter,  let  us  apply  them  to  the  imaginary  case  of  a  newly  dis- 
covered element.  The  bromide  of  the  element  is  found  to  be 
easy  of  preparation  and  to  be  volatile.  The  bromide  contains 
30  per  cent  of  the  element  (and  therefore  70  per  cent  of  bromine), 
and  its  vapor  density  referred  to  air  is  11.8.  The  analysis  can 
always  be  made  much  more  accurately  than  the  measurement  of 
vapor  density,  so  that  the  former  number  is  more  trustworthy 
than  the  latter. 

To  find  the  equivalent  of  the  element,  that  is,  the  amount  com- 
bined with  79.92  parts  (the  atomic  weight)  of  bromine,  we  have  the 
proportion  70  :  30  ::  79.92  :  x,  from  which  x  =  34.3.  The  atomic 
weight  must  be  this,  or  some  small  multiple  of  it. 

The  G.M.V.  of  air  weighs  28.955  g.  (p.  101).  Hence  the  same 
volume  of  the  vapor  of  this  bromide,  which  is  1 1 .8  times  as  heavy  as 
air,  will  weigh  28.955  X  11.8,  or  341.67  g.  This  is  therefore  the 
molar  weight  of  the  compound. 

Now  30  per  cent  of  this  is  the  new  element: 

341.67  X  30  -^  100  =  102.5. 

Now  34.3  parts  of  the  element  combined  with  79.92  parts'of  bro- 
mine. Evidently  the  atomic  weight  of  the  element  is  3  X  34.3  = 
102.9,  the  difference  being  due  to  error  in  determining  the  density. 
So  long  as  no  other  volatile  compound  is  known,  we  adopt  this  as 
the  atomic  weight.  The  rest  of  the  molar  weight  (239  parts  = 
3  X  79.92)  is  bromine.  Thus  the  formula  of  the  compound  is 
ElBr3,  and  from  this  we  see  that  the  element  is  trivalent. 

In  case  no  volatile  compound  of  the  element  can  be  formed,  the 
weight  combining  with  79.92  parts  of  bromine  is  measured  as  before. 
Then  some  of  the  free  simple  substance  is  made,  say  by  electrol- 
ysis, and  its  specific  heat  is  determined.  The  sp.  ht.  is  about 
0.063.  Application  of  Dulong  and  Petit's  law  then  gives  the 
atomic  weight.  The  product  34.3  X  0.063  is  equal  to*2.161. 
Hence,  the  equivalent  must  be  multiplied  by  3  to  give  the  atomic 


APPLICATIONS   OF   MOLECULAR   EQUATIONS  119 

weight,  for  this  raises  the  product  to  6.48,  which  is  within 
the  limits.  Thus  the  value  of  the  atomic  weight  is  102.9,  as 
before. 

Replies  to  Questions  about  Difficulties.  —  The  beginner 
always  becomes  confused  over  one  or  more  of  the  points  raised  by 
the  following  questions: 

1.  Why  was  32  g.  of  oxygen  taken  as  the  standard  for  molecular 
weights,  rather  than  16  g.?     Read  p.  107  and  footnote  to  p.  104. 

2.  If  O2  is  the  smallest  mass  of  oxygen,  why  do  we  have  formulae 
like  H20  and  HC1O?     02  is  the  smallest  mass  of  free  oxygen,  but 
in  combination  half  as  much  occurs  in  many  molecules.     Read 
pp.  105,  110,  and  111. 

3.  Why  is  not  the  atomic  weight  of  an  element  ascertained  by 
simply  measuring  the  density  of  the  elementary  substance?     Read 
pp.  Ill,  last  par.,  and  117,  second  par. 

4.  Can  we  not  deduce  the  valence  of  an  element  from  knowing 
the  number  of  atoms  in  its  molecules,  and  vice  versa?    Some  molec- 
ular formulae  and  valences  are :  H2X,  O2n,  CV,  Znn,  also  Hg  (uni- 
valent  and  bivalent),  P4  (trivalent  and  quinquivalent),  and  Ss 
(bivalent  and  sexivalent).     There  is  no  relation,  either  observable 
or  to  be  expected. 

5.  Do  the  molecular  weights,  oxygen  =  32  and  hydrogen  =  2, 
mean  that  the  molecules  of  oxygen  are  larger  than  are  those  of 
hydrogen?     This  is  the  ratio  of  their  weights,  but  none  of  the 
phenomena  discussed  in  this  chapter  are  influenced  appreciably 
by  their  relative  sizes,  and  therefore  none  of  them  give  any  in- 
formation on  the  subject.     Read  the  footnote  to  p.  102. 

Exercises.  — 3L    The  weight  of  1 1.  of  a  gas  at  0°  and  760  mm.  is 

5.236  g.  What  is  the  density  referred  (a)  to  air  (air  =  1)  and  (6) 
to  hydrogen,  and  (c)  what  is  the  molecular  weight  (pp.  101,  102)? 

2.  The  density  of  a  gas,  referred  to  air,  is  6.7.     What  is  the 
weight  of  1  1.  (p.  101),  and  what  is  the  molecular  weight  (p.  118)? 

3.  The  molecular  weight  of  a  substance  is  65.     What  is  the 
density  referred  to  air,  and  what  is  the  weight  of  1  1.? 

4.  The  chloride  of  a  new  element  contains  38.11  per  cent  of 
chlorine  and  61.89  per  cent  of  the  element.     The  vapor  density  of 
the  compound  referred  to  air  is  12.85.     What  is  the  atomic  weight 


120  COLLEGE   CHEMISTKY 

of  the  element,  so  far  as  investigation  of  this  one  substance  can 
give  it  (p.  118)?    What  is  its  valence? 

5.  If  the  molecular  weight  of  oxygen  were  taken  as  100,  what 
would  be  the  volume  of  the  G.M.V.  (p.  101)?    What,  on  the  same 
scale,  would  be  the  molecular  weight  of  water,  and  what  would  be 
the  atomic  weights  of  hydrogen  and  chlorine  (pp.  101,  105)? 

6.  In  future  nothing  but  molecular  formulae  of  free  elements 
must  be  used  (p.  111).    Write  in  molecular  form  ten  of  the  equa- 
tions involving  gases  which  are  found  in  the  preceding  chapters. 

7.  If  a  new  form  of  oxygen  were  found,  such  that  one  volume 
of  it  required  four  volumes  of  hydrogen  to  produce  water,  what 
would  be  its  molecular  formula  (p.  114)?     What  would  be  the 
weight  of  22.4  1.? 


CHAPTER  X 
SOLUTION 

SOLUTIONS  are  so  constantly  used  in  chemistry  that  some 
knowledge  of  their  properties  is  desirable  in  order  that  we  may 
employ  them  intelligently.  In  what  follows,  we  give  a  preliminary 
account  of  some  of  the  simpler  facts  about  solution. 

General  Properties  of  Solutions.  —  A  solid  may  be  dis- 
tributed through  a  liquid,  either  by  being  simply  suspended  (p.  12) 
in  the  latter  (mixture),  or  by  being  dissolved  in  it  (solution). 
Similarly  a  liquid  may  be  suspended  in  droplets  in  another  liquid 
(emulsion),  as  in  milk,  or  it  may  be  dissolved.  It  is  usually  easy  to 
distinguish  between  the  two  cases,  for  a  suspended  substance  settles 
or  separates  sooner  or  later  (like  the  fats  in  milk  —  as  cream), 
while  a  dissolved  substance  shows  no  such  tendency.  The  cases  are 
exceptional  where  the  subdivision  of  a  suspended  substance  is  so 
minute  (colloidal  suspension,  q.v.),  as  to  make  its  retention  by  filter 
paper  impossible.  If  a  liquid  is  opalescent  or  opaque,  then  we  have 
a  case  of  suspension.  A  solution  is  a  clear,  transparent,  perfectly 
homogeneous  liquid,  in  which  the  dissolved  substance  seems  to 
have  been  dispersed  so  completely  that  the  liquid  cannot  be  dis- 
tinguished by  the  eye  from  a  pure  substance. 

There  is  no  limit  to  the  amount  of  dissipation  which  may  thus 
be  produced.  A  single  fragment  of  potassium  permanganate,  for 
example,  which  gives  a  very  deep  purple  solution  in  water,  may  be 
dissolved  in  a  liter  or  even  in  twenty  liters  of  water,  and  the  purple 
tinge  which  it  gives  to  the  liquid  will  still  be  perfectly  perceptible 
in  every  part  of  the  larger  volume.  The  qualitative  characteristics, 
therefore,  of  solution  are  absence  of  settling,  homogeneity,  and  ex- 
tremely minute  subdivision  of  the  dissolved  substance. 

The  Scope  of  the  Word.  —  The  word  solution  is  used  for  other 
systems  than  those  containing  a  solid  body  dissolved  in  a  liquid. 

121 


122  COLLEGE    CHEMISTRY 

Thus,  liquids  also  may  be  dissolved  in  liquids,  as  alcohol  in  water. 
Again,  if  we  warm  ordinary  water,  bubbles  of  gas  appear  on  the 
sides  of  the  vessel  before  the  water  has  approached  the  boiling- 
point.  They  are  found  to  be  gas  derived  from  the  air.  Agitation 
of  any  gas  with  water  results  in  the  solution  of  a  large  or  small 
quantity  of  the  gas,  and  heat  will  usually  drive  the  gas  out  again. 
It  appears  therefore  that  solids,  liquids,  and  gases  can  equally  form 
solutions  in  liquids. 

The  absorption  of  hydrogen  by  palladium  (at  all  events  after  a 
certain  point),  and  by  iron,  takes  place  in  accordance  with  the 
same  laws  as  the  solution  of  solids  in  liquids,  and  the  results  may  be 
described  therefore  as  true  solutions.  Liquids  are  in  some  cases 
absorbed  by  solids,  and  homogeneous  mixtures  of  solids  with  solids 
are  perfectly  familiar.  The  sapphire  is  a  solution  of  a  small 
amount  of  a  strongly  colored  substance,  in  a  large  amount  of  color- 
less aluminium  oxide.  It  may  therefore  be  stated  that  solution  of 
gases,  liquids,  and  solids  in  solids  appears  to  be  possible. 

Limits  of  Solubility.  —  The  next  question  which  naturally 
occurs  to  us  is  as  to  whether  the  mingling  of  two  substances  in  this 
manner  has  any  limits.  We  find  that  the  results  of  experiment  in 
this  direction  may  be  divided  into  two  classes.  Some  pairs,  of 
liquids  particularly,  may  be  mixed  in  any  proportions  whatever. 
Alcohol  and  water,  and  glycerine  and  water  are  such  pairs.  On 
the  other  hand,  at  the  ordinary  laboratory  temperature,  we  can 
scarcely  take  a  fragment  of  marble  (CaCOa)  so  small  that  it  will 
dissolve  completely  in  100  c.c.  of  pure  water,  for  only  0.00013  g. 
dissolves.  Under  the  same  conditions  any  amount  of  potassium 
chlorate  up  to  about  5  g.  will  completely  disappear  after  vigorous 
stirring,  while  90  g.  of  ordinary  Epsom  salts  (hydrated  magnesium 
sulphate),  but  not  more,  may  be  dissolved  in  about  the  same 
amount  of  water.  In  fact,  most  solids  may  be"  dissolved  in  a  liquid 
only  up  to  a  certain  limit,  which  with  different  solids  may  range 
from  a  scarcely  perceptible  to  a  very  large  amount.  No  substance 
is  absolutely  insoluble.  But  for  the  sake  of  brevity  we  call  marble, 
for  example,  "insoluble"  because  in  most  connections  it  may  be  so 
considered. 

Chemists  have  not  yet  succeeded  in  explaining  these  differences 
in  solubility,  which  are  often  so  surprising.  Thus,  guncotton  is 


SOLUTION  123 

soluble  in  a  mixture  of  alcohol  and  ether,  but  not  in  these  liquids 
separately,  while  cellulose  acetate  (an  allied  substance,  used  in 
making  artificial  horse-hair)  is  soluble  in  these  liquids  separately, 
but  not  in  the  mixture. 

Recognition  and  Measurement  of  Solubility.  —  The  only 
method  of  recognizing  with  certainty  whether  a  solid  is  soluble  in  a 
liquid  or  not  is  to  filter  the  mixture  and  evaporate  a  few  drops  of 
the  filtrate  on  a  clean  watch-glass.  For  learning  how  much  of  the 
body  is  contained  in  a  given  solution,  a  weighed  quantity  of  the 
solution  is  evaporated  to  dryness  and  the  weight  of  the  residue 
determined.  i 

It  must  be  stated  explicitly  that  in  going  into  solution,  as  we 
have  used  the  term,  a  compound  dissolves  as  a  whole  and,  if  the 
compound  is  pure  (p.  4),  any  residue  has  the  same  chemical  com- 
position as  the  part  which  has  dissolved.  If  the  residue  is  a 
different  substance,  a  chemical  interaction  with  the  solvent  has 
occurred.  If,  on  evaporation,  a  different  substance  remains,  there 
has  also  been  chemical  action. 

Terminology.  —  In  order  to  describe  the  relations  of  the  com- 
ponents of  a  solution,  certain  conceptions  and  corresponding 
technical  expressions  are  required. 

It  is  customary  to  speak  of  the  substance  which,  like  water  in 
most  cases,  forms  the  bulk  of  the  solution,  as  the  solvent.  To 
express  the  substance  which  is  dissolved,  the  word  solute  is  fre- 
quently used,  and  will  be  employed  when  we  wish  to  avoid  circum- 
locution. 

The  amount  of  the  substance  which  has  been  dissolved  by  a 
given  quantity  of  the  solvent  is  described  as  the  concentration  of 
the  solution.  A  solution  containing  a  small  proportion  of  the 
dissolved  body  is  called  dilute;  it  has  a  small  concentration. 
One  which  contains  a  larger  amount  is  more  concentrated.  Very 
"strong"  solutions  are  frequently  spoken  of  simply  as  concentrated 
solutions.  The  partial  removal  of  the  solvent  (as  by  evaporation) 
is  called  concentrating,  its  total  removal  evaporating  to  dryness. 

Finally,  since  there  is  a  limit  to  the  solubility  of  most  substances, 
a  solution  is  described  as  saturated  when  the  solute  has  given  as 
much  material  to  the  solvent  as  it  can.  This  state  is  reached  after 


124  COLLEGE    CHEMISTRY 

prolonged  agitation  with  an  excess  of  the  gas,  of  the  liquid,  or  of 
the  finely  powdered  solid,  as  the  case  may  be  (see  pp.  127,  133). 
Other  things  being  equal,  the  larger  the  excess,  the  sooner  satura- 
tion is  attained.  The  maximum  concentration  attainable  in  this 
way  is  called  the  solubility  of  the  substance  in  a  given  solvent. 
Note  that  a  saturated  solution  need  not  also  be  a  concentrated  one. 
It  will  be  very  dilute,  if  the  solute  is  but  slightly  soluble. 

Units  Used  in  Expressing  Concentrations.  —  The  concen- 
trations of  solutions,  saturated  and  otherwise,  are  sometimes 
expressed  in  physical,  and  sometimes  in  chemical,  units  of  weight. 
When  physical  units  are  employed,  we  give  the  number  of  grams 
of  the  solute  held  in  solution  by  one  hundred  grams  of  the  solvent. 

The  solubilities  at  18°  of  one  hundred  and  forty-two  bases  and 
salts  are  given  in  a  table  printed  inside  the  cover,  at  the  front  of 
this  book. 

When  chemical  units  of  weight  are  employed,  two  different  plans 
are  possible,  and  both  are  in  use.  Either  the  equivalent  (p.  65)  or 
the  molecular  weights  may  be  taken  as  a  basis  of  measurement. 
In  the  former  case,  the  solutions  are  called  normal  solutions,  and  in 
the  latter,  molar  solutions. 

A  normal  solution  contains  one  gram-equivalent  of  the  solute  in  one 
liter  of  solution  (not  in  1  1.  of  solvent).  The  word  " equivalent" 
has  been  used  hitherto  only  of  elements,  and  this  application  of  the 
expression  involves  an  extension  of  its  meaning.  An  equivalent 
weight  of  a  compound  is  that  amount  of  it  which  will  interact  with 
one  equivalent  of  an  element.  Thus,  a  formula-weight  of  hydro- 
chloric acid  HC1  (36.5  g.)  is  also  an  equivalent  weight,  for  it  con- 
tains 1  g.  of  hydrogen,  and  this  amount  of  hydrogen  is  displaceable 
by  one  equivalent  weight  of  a  metal.  A  formula-weight  of 
sulphuric  acid  H2S04  (98  g.),  however,  contains  two  equivalents 
of  the  compound,  and  a  formula-weight  of  aluminium  chloride 
A1C13  (133.5  g.)  three  equivalents.  Hence  normal  solutions  of 
these  three  substances  contain  respectively  36.5  g.  HC1,  49  g. 
H2S04,  and  44.5  g.  A1C13  per  liter  of  solution.  The  special 
property  of  normal  solutions  is,  obviously,  that  equal  volumes  of 
two  of  them  contain  the  exact  proportions  of  the  solutes  which  are 
required  for  complete  interaction.  Solutions  of  this  kind  are  much 
used  in  quantitative  analysis.  We  frequently  use  also  decinormal 


SOLUTION  125 

or  one-tenth  normal  solutions  (0.1  N  or  JV/10),  and  seminormal 
(0.5  N  or  2V/2),  and  six  times  normal  solutions  (6  N),  and  so 
forth. 

A  molar  solution  contains  one  mole  (gram-molecular  weight)  of  the 
solute  in  one  liter  of  solution  (not  in  1  1.  of  solvent).  When  molec- 
ular formulae  (p.  109)  are  used,  this  means  one  gram-formula 
weight  per  liter.  In  the  cases  cited  above,  the  molar  solution 
contains  36.5  g.  HC1,  98  g.  H2SO4,  and  133.5  g.  A1C13  per  liter. 
As  will  be  seen,  the  concentrations  of  molar  and  normal  solutions 
are  necessarily  identical  when  the  radicals  are  univalent. 

Solution  One  of  the  Physical  States  of  Aggregation  of 
Matter.  —  When  a  solid  body  dissolves  in  a  liquid,  the  properties 
of  the  body  undergo  a  very  marked  change,  which  to  all  appearance 
might  be  chemical.  Yet,  besides  the  ease  with  which  a  liquid  may 
be  removed  by  evaporation  and  the  solid  recovered  unchanged,  we 
note  particularly  that  the  concentration  of  a  saturated  solution 
cannot  be  expressed  in  terms  of  integral  multiples  of  the  atomic 
weights.  We  shall  see  also  that  the  quantity  of  a  solid  which  a 
liquid  may  take  up  varies  with  the  slightest  change  in  temperature. 
Now  we  do  not  find  the  composition  of  chemical  compounds  so  to 
vary.  The  solution  of  a  solid  may  therefore,  in  general,  be  likened 
to  a  change  in  state  of  aggregation,  similar  to  the  conversion  of  a 
liquid  into  a  gas  or  a  solid  (see  p.  126). 

As  in  other  changes  of  state,  so  in  the  process  of  solution,  heat  is 
always  liberated  or  absorbed.  This  is  known  as  heat  of  solution. 
Thus,  one  formula-weight  of  sulphuric  acid,  in  dissolving  in  a  large 
volume  of  water,  liberates  39,170  calories,  and  one  formula-weight 
of  ammonium  chloride,  in  dissolving,  absorbs  3880  calories. 

As  there  is  danger  of  confusion  arising,  we  may  repeat  that  a 
compound  is  homogeneous  and  its  composition  is  expressible  in 
chemical  units  of  weight;  a  saturated  solution  is  homogeneous  but 
its  concentration  varies  with  temperature  so  that  atomic  weights 
cannot  be  used  to  describe  its  composition;  a  mixture  of  two  solids, 
or  an  emulsion  of  two  liquids,  is  neither  homogeneous  nor  in  any 
way  definite  in  composition. 

Molecular  View  of  the  State  of  Solution.  —  Accepting  solu- 
tion as  a  physical  state  of  aggregation,  we  may  now  apply  the  same 


126  COLLEGE    CHEMISTRY 

molecular  conceptions  to  the  explanation  of  the  behavior  of  a  sub- 
stance in  solution  as  to  matter  in  the  gaseous  or  liquid  states.  We 
saw  that  a  solid  body,  which  is  ordinarily  condensed  in  a  small 
space,  can  be  disseminated  by  the  use  of  a  solvent  through  a  very 
large  one.  The  molecules  of  the  solid  become  scattered  like  those 
of  a  gas  or  vapor  through  a  much  greater  volume.  We  may  re- 
gard the  dissolved  substance  as  being,  practically,  in  a  gaseous  or 
quasi-gaseous  condition.  The  molecules  are  torn  apart  from  one 
another,  their  cohesion  is  overcome,  and  their  freedom  of  motion 
is  in  a  measure  restored.  It  is  true  that  they  could  not  continue 
to  occupy  this  large  volume  for  a  moment  in  the  absence  of  the 
solvent.  But  we  may  bring  this  into  relation  with  the  case  of  a 
vapor  by  saying  that  a  solid  body,  like  common  salt,  can  evapo- 
rate (i.e.,  "dissolve")  at  the  ordinary  temperature,  and  occupy  a 
large  space,  only  when  that  space  is  already  filled  with  a  suitable 
liquid.  The  latter  acts  as  a  vehicle  for  the  particles  of  the  solid. 
A  volatile  liquid,  on  the  contrary,  can  dissolve  in  an  empty  space 
and  fill  it  with  its  particles  without  any  vehicle  being  required. 

This  conception  of  the  quasi-gaseous  condition  of  a  dissolved  sub- 
stance would  be  simply  fantastic  if  it  did  not  lead  us  to  a  better 
understanding  of  the  behavior  of  solutions.  It  does  successfully 
explain  many  things,  such  as  diffusion,  osmotic  pressure,  and  satu- 
ration (see  next  section). 

It  is  easy  to  show  that,  if  we  place  a  quantity  of  the  pure  solvent 
(Fig.  55)  above  a  concentrated  solution  of  a  substance,  and  then  set 
the  arrangement  aside,  the  dissolved  body  slowly  makes  its  way 
through  the  liquid  (Fig.  56),  obliterating  the  original  plane  of  sepa- 
ration. Eventually  the  dissolved  body  scatters  itself  uniformly 
through  the  whole.  In  other  words,  the  particles  of  the  dissolved 
substance  exhibit  the  property  of  diffusion  in  the  same  way  as  do 
those  of  gases. 

When  the  diffusion  of  a  gas  is  resisted  by  a  suitable  partition,  we 
find  that  pressure  is  exercised  upon  the  walls  of  the  vessel  and  upon 
the  partition.  It  is  possible  to  show  that  the  particles  of  a  dis- 
solved substance  exercise  a  pressure  of  a  very  similar  kind.  This 
pressure  is  spoken  of  as  osmotic  pressure.  This  pressure  is  found 
to  be  proportional  to  the  concentration  of  the  solution,  just  as 
gaseous  pressure  is  proportional  to  the  concentration  of  the  gas 
(Boyle's  law). 


SOLUTION 


127 


Molecular   View  of  the  Process  of  Solution.  —  We  may 

now  apply  the  same  ideas  to  the  process  of  dissolving,  with  a  view 


FIG.  55. 


FIG.  56. 


more  especially  to  explaining  why  this  process  ceases,  in  spite  of  the 
presence  of  excess  of  the  solute,  when  a  certain 
concentration  has  been  reached.     If  some  of 
the  material  dissolves,  why  not  more? 

Let  us  suppose  that  it  is  the  dissolving  of 
common  salt  in  water  (Fig.  57)  which  we  wish 
to  explain  in  detail.  We  believe  that  in  the 
solid  substance  the  molecules  are  closely 
packed  together,  while  in  the  solution  they 
are  rather  sparsely  distributed.  The  process 
of  solution  must  consist  in  the  loosening  of  the 
molecules  on  the  surface  and  their  passage  into 
the  liquid.  By  diffusion,  the  free  molecules 
will  gradually  move  away  from  the  neighbor- 
hood of  the  surface  of  the  solid  and  make  room 
for  others,  and  thus,  if  the  system  remains 
undisturbed,  the  liquid  will  eventually  become 
a  solution  of  uniform  concentration.  If  a  large 
enough  amount  of  the  solid  has  been  provided,  the  ultimate  condi- 
tion will  be  that  of  a  saturated  solution  with  excess  of  the  solid 


FIG.  57. 


128  COLLEGE    CHEMISTRY 

beneath.  If  we  had  proper  means  of  measuring  it,  the  tendency 
of  the  molecules  to  leave  the  solid  in  the  presence  of  a  given 
liquid  would  give  the  effect  of  a  kind  of  pressure.  This  is  spoken 
of  as  solution  pressure. 

Now  the  molecules,  after  having  entered  the  liquid,  move  in 
every  direction,  and  consequently  some  of  them  will  return  to  the 
solid  and  attach  themselves  to  it.  The  frequency  with  which  this 
will  occur  will  be  greater  as  the  crowding  of  particles  hi  the  liquid 
increases,  so  that  a  stage  will  eventually  be  reached  at  which  the 
number  of  molecules  leaving  the  solid  will  be  no  greater  than  that 
landing  upon  it  in  a  given  time.  If  the  whole  of  the  liquid  has 
meanwhile  become  equally  charged  with  dissolved  molecules,  there 
will  be  no  chance  that  the  field  of  liquid  immediately  round  the  solid 
will  lose  them  by  diffusion,  so  that  a  condition  of  balance  or  equi- 
librium (p.  89)  will  have  been  established:  NaCl  (solid)  +±  NaCl 
(dslvd).  The  motion  of  the  particles  in  the  liquid  produces  what 
we  have  called  osmotic  pressure;  and  when  the  osmotic  pressure,  by 
the  continual  increase  in  the  number  of  dissolved  molecules,  becomes 
equal  to  the  solution  pressure,  increase  in  concentration  of  the  solu- 
tion ceases.  It  is  at  this  point  that  we  speak  of  the  solution  as  being 
saturated  with  respect  to  the  particular  substance  dissolving.  The 
analogy  to  vapor  tension  and  vapor  pressure  (p.  88)  is  evident. 
The  foregoing  explanation  should  be  compared  carefully  with  that 
given  in  the  section  on  the  molecular  relations  in  liquids,  and  in 
that  on  equilibrium  (pp.  81,  89-90). 

Conditions  Affecting  the  Solubility  of  a  Gas.  —  When  the 
dissolving  substance  is  a  gas,  led  through,  or  confined  above  the 
liquid  at  a  definite  pressure,  the  gas  dissolves  until  a  state  of  equi- 
librium between  dissolving  and  emission  is  reached,  for  example, 
Oxygen  (gas)  <F±  Oxygen  (dslvd),  and  the  liquid  is  then  saturated 
with  the  gas. 

It  is  found,  as  the  molecular  theory  would  lead  us  to  expect, 
that  the  concentration  of  the  saturated  solution  of  a  gas  is  propor- 
tional to  the  pressure  at  which  the  gas  is  supplied  (Henry's  law). 

This  equilibrium,  Gas  (gaseous)  <±  Gas  (dslvd),  can  be  reached, 
naturally,  from  the  other  direction,  namely  by  starting  with  a 
solution  of  the  gas  and  a  space  above  the  solution  containing,  at 
first,  none  of  the  gas.  The  gas  leaves  the  solution  until  the  rates 


SOLUTION  129 

of  emission  and  return  become  equal.  Hence,  a  gas  may  be  en- 
tirely removed  from  solution  by  bubbling  a  foreign  gas  through 
the  liquid.  The  bubbles  furnish  the  space  to  receive  the  emitted 
gas,  and  have  a  large  surface,  so  that  the  process  goes  on  rapidly. 
The  bubbles  also  escape,  and  carry  with  them  the  emitted  gas,  so 
that,  in  this  case,  there  is  no  re-solution.  This  is  a  case  of  nullify- 
ing one  of  the  two  opposed  tendencies  (p.  90). 

When  a  mixture  of  two  gases  is  shaken  with  a  liquid,  the 
gases  behave  independently  of  each  other  (Dalton's  law,  p.  72). 
Each  has  the  same  pressure,  and  therefore  the  same  solubility, 
as  it  would  possess  if  it  alone  occupied  the  whole  space  above  the 
liquid. 

Two  Immiscible  Solvents:  Law  of  Partition.  —  An  interest- 
ing application  of  the  same  ideas  may  be  made  to  a  case  which 
occurs  very  commonly  in  chemical  work.  If  we  shake  up  a  small 
particle  of  iodine  with  water,  we  find  that  it  dissolves  slowly,  giving 
eventually  a  saturated  but  very  dilute  solution.  If  now  ether  in 
sufficient  quantity  be  shaken  with  the  aqueous  solution,  the  greater 
part  of  the  iodine  will  find  its  way  into  the  ether,  and  be  contained 
in  the  brown  layer  which  rises  to  the  top.  The  process  of  re- 
moving a  substance  partially  from  solution  in  one  solvent  and 
securing  it  in  another  is  called  extraction.  We  find  in  such  cases 
that  neither  solvent  can  entirely  deprive  the  other  of  the  whole  of 
the  dissolved  substance,  if  the  latter  is  soluble  in  both  independ- 
ently. A  state  of  equilibrium  is  finally  reached:  I2  (in  Aq)  +±I% 
(in  ether).  The  partition  of  the  substance  takes  place  in  propor- 
tion to  its  solubility  in  each  solvent.  It  is  found  that  any  amount 
of  the  solute,  up  to  the  maximum  the  system  can  contain,  provided 
this  does  not  involve  too  high  a  concentration  in  either  solvent,  is 
divided  so  that  the  ratio  of  the  concentrations  in  the  two  solvents 
is  always  the  same.  In  the  case  of  iodine  divided  between  water 
and  ether,  this  ratio  is  about  1  :  200. 

This  principle  is  used  in  Parke's  process  (q.v.)  for  extracting 
silver  from  molten  lead,  by  means  of  melted  zinc  as  the  second 
solvent.  It  is  employed  in  separating  interesting  compounds  from 
animal  secretions  and  vegetable  extracts,  and  in  purifying  such 
compounds.  Nicotine  from  tobacco  and  cocaine  from  coca- 
leaves,  are  secured  in  this  way. 


130  COLLEGE    CHEMISTRY 

Influence  of  Temperature  on  Solubility.  —  The  quantity  of 
a  substance  which  we  can  dissolve  in  a  fixed  amount  of  a  given 
solvent  depends  very  largely  upon  the  temperature  of  both.  Usu- 
ally the  solubility  increases  with  rise  in  temperature.  Measure- 
ments may  be  made  by  the  method  described  before  (p.  123),  using 
excess  of  the  finely  powdered  solute  with  different  portions  of  the 
same  solvent  in  vessels  kept  at  different  temperatures.  The  most 
useful  way  of  representing  the  results  is  to  plot  them  graphically. 
The  diagram  (Fig.  58)  shows  the  curves  for  a  few  familiar  sub- 
stances. The  ordinates  represent  the  number  of  grams  of  the 
anhydrous  compound  which  is  held  in  solution  by  100  g.  of  water 
in  each  case.  The  abscissae  represent  the  temperatures.  The  con- 
centration for  any  temperature  can  be  read  off  at  once.  Thus,  100  g. 
of  water  holds  13  g.  of  potassium  nitrate  in  solution  at  0°  and  150  g. 
at  73°.  The  increase  in  solubility  is  here  enormous.  On  the 
other  hand,  the  same  quantity  of  water  will  hold  35.6  g.  of  sodium 
chloride  in  solution  at  0°  and  39  g.  at  100°.  The  difference  is  shown 
at  once  when  we  examine  the  curves  and  observe  that  the  line  repre- 
senting the  solubility  of  sodium  chloride  scarcely  rises  at  all  between 
0°  and  100°,  while  that  of  potassium  nitrate  is  extremely  steep. 

Cases  in  which  the  solubility  decreases  with  rise  in  temperature 
are  less  common.  The  solubility  of  slaked  lime  (calcium  hydroxide 
Ca(OH)2,  used  to  make  limewater)  is  0.175  g.  at  20°  and  0.078  g. 
at  100°.  Anhydrous  sodium  sulphate  Na2SO4  (Fig.  59,  p.  132)  is 
another  illustration. 

Equilibrium  in  a  Saturated  Solution.  —  Once  a  solution  has 
become  saturated,  the  dissolving  substance  remains  thereafter  un- 
changed in  amount.  A  greater  excess  of  the  solute  forces  no  more 
matter  into  solution  than  does  a  small  excess. 

It  should  be  clearly  understood,  however,  that  an  exchange  of 
molecules  (p.  128)  is  still  going  on  between  the  solid  and  the  solu- 
tion. That  this  conception  is  correct  may  be  shown  in  various 
ways.  Thus,  if  a  crystal,  the  edges  or  corners  of  which  have  been 
broken,  is  suspended  in  a  saturated  solution  of  the  same  substance, 
it  neither  increases  nor  diminishes  in  weight.  Yet  we  find  that  the 
imperfections  are  removed,  and  that  this  takes  place  by  the  solu- 
tion of  a  portion  of  the  substance  from  the  perfect  surfaces  and  its 
deposition  upon  the  imperfect  ones. 


SOLUTION 


131 


132 


COLLEGE    CHEMISTRY 


Supersaturated  Solutions.  —  Another  very  striking  proof  of 
this  may  be  obtained  by  saturating  water  with  ordinary  Glauber's 
salt  (hydrated  sodium  sulphate  Na2S04,10H2O)  at,  say,  30°,  at 
which  temperature  100  c.c.  of  water  hold  in  solution  40  g.  of 


o° 


40°      50° 
Temperature 
Fia.  59. 


90°    100° 


Na2SO4  (Fig.  59).  The  excess  of  the  solid  is  carefully  and  com- 
pletely separated  from  the  liquid,  and  the  latter  is  allowed  to  cool, 
say  to  15°,  in  a  flask  loosely  stoppered  with  cotton. 
The  solution  now  contains  a  much  larger  amount 
of  sodium  sulphate  (N^SC^)  than  at  its  present 
temperature  it  could  acquire  from  contact  with 
Glauber's  salt  (13  g.  at  15°).  Yet  in  the  absence 
of  a  crystal,  with  which  the  above  described  ex- 
change could  take  place,  no  deposition  of  the  dis- 
solved substance  begins.  The  solution  may  be 
kept  indefinitely  without  alteration.  The  intro- 
duction, however,  of  the  minutest  fragment  of 
the  decahydrate  at  once  starts  the  exchange,  and 
this  is  necessarily  very  much  to  the  disadvantage 
of  the  solution  and  the  advantage  of  the  crystal : 
10H2O  +  Na2SO4  (dslvd)  ^±  Na2S04,10H20  (solid).  The  latter 
therefore  forms  the  center  of  a  radiating  mass  of  blade-like  proc- 
esses, which  sprout  with  astonishing  rapidity  through  the  liquid 
(Fig.  60). 


FIG.  60. 


SOLUTION  133 

Usually  the  cooling  of  a  concentrated  solution  leads  to  the  almost 
immediate  appearance  of  crystals  spontaneously,  and  the  substance 
is  deposited  gradually  as  the  temperature  falls.  But  solutions 
of  a  number  of  common  substances,  such  as  sodium  thiosulphate 
(photographer's  "hypo")  and  sodium  chlorate,  behave  like  that 
of  sodium  sulphate.  They  are  said  to  have  a  tendency  to  give 
supersaturated  solutions.  In  general,  crystallization  can  be  started 
only  by  introduction  of  a  specimen  of  the  same  substance.  The 
smallest  particle  of  the  right  material  floating  in  the  air,  if  it  gams 
admission,  will  bring  about  the  result.  This  shows  the  importance 
of  the  interchange  of  molecules,  of  which  we  have  spoken,  for 
establishing  equilibrium. 

This  phenomenon  is  similar  to  the  supercooling  of  water  (p.  86), 
which  results  in  crystallization  (freezing)  when  a  fragment  of  ice 
is  dropped  in. 

Definition  of  a  Saturated  Solution:  A  Warning.  —  To  avoid 
a  common  misconception,  it  must  be  noted  that  a  saturated  solution 
must  not  be  denned  as  one  containing  all  of  the  solute  that  it  can  hold. 
A  supersaturated  solution  holds  more.  The  saturated  solution  is 
one  which  contains  all  of  the  dissolved  solute  that  it  can  acquire  from 
the  undissolved  solute.  Better  still,  it  is  that  solution  which,  when 
placed  with  excess  of  the  solute,  is  found  to  be  in  equilibrium. 

It  must  be  clearly  understood  that  solution  is  not  a  process  of 
filling  the  pores  of  the  liquid.  If  that  were  true,  approximately 
equal  weights  of  all  substances  would  find  accommodation  in  equal 
volumes  of  water.  The  fact  is  that,  for  example,  100  c.c.  of  water 
can  dissolve  195  g.  of  silver  fluoride,  but  only  0.00000035  g.  of 
silver  iodide,  although  the  space  available  in  the  solvent  (if  there 
is  any  free  space)  is  the  same  in  both  cases. 

The  same  conclusion  is  reached  when  we  consider  that  two  forms 
of  the  same  compound  have  different  solubilities.  Thus,  at  20°, 
Na2SO4,10H20  can  give  about  18  g.  of  Na^SCX  to  100  c.c.  of  water 
(Fig.  59).  But  anhydrous  sodium  sulphate  Na^SC^  at  20°  gives 
59  g.  to  the  same  amount  of  water  (read  up  to  dotted  line,  Fig.  59). 
Note  that  in  the  diagram  (Fig.  59)  the  solubility  curve  of  Na2S04, 
lOHoO  comes  to  an  end  at  32.4°.  At  this  temperature  the  solid 
melts  and  decomposes,  so  that  measurements  with  this  solid  be- 
yond that  temperature  are  impossible. 


134  COLLEGE    CHEMISTRY 

Influences  of  the  Solute  Upon  the  Solvent.  —  These  in- 
fluences are  of  two  classes.  In  one  of  these  classes,  equal  num- 
bers of  dissolved  molecules  of  different  substances  produce  the 
same  amount  of  change.  The  effect  appears,  therefore,  to  be 
largely  due  to  mechanical  causes.  Of  this  nature  are  the  lowering 
of  the  freezing-point  of  the  liquid,  the  lowering  of  its  vapor  tension 
and  the  raising  of  its  boiling-point,  and  the  value  of  the  osmotic 
pressure  (see  below). 

In  the  other  class,  the  effect  varies  with  the  substance  dissolved. 
The  changes  in  volume  (see  below)  belong  to  this  class. 

Freezing-Points  of  Solutions:    Freezing  Mixtures.  —  The 

freezing  of  a  dilute  solution  consists,  usually,  in  the  crystalliza- 
tion of  some  of  the  pure  solvent  only.  The  presence  of  a  dissolved 
body  tends  to  prevent  this  freezing,  and  so  solutions  can  be  frozen 
only  at  temperatures  below  those  at  which  the  pure  solvent  would 
freeze.  Thus,  one  gram-molecular  weight  of  any*  substance, 
such  as  sugar  (342  g.)  or  alcohol  (46  g.),  dissolved  in  1000  c.c. 
(1  liter)  of  water,  will  cause  the  water  to  freeze  at  —1.86°  instead 
of  at  0°. 

This  explains  why  sea  water  is  much  less  often  frozen  in  cold 
weather  than  is  fresh  water.  It  should  be  noted,  also,  that  the  ice 
formed  in  salt  water  is  free  from  salt. 

This  fact  likewise  explains  why  salt  thrown  on  ice  causes  the 
latter  to  melt.  A  saturated  solution  of  salt  does  not  freeze  until 
cooled  to  —21°  (  —  6°  F.),  and  it  then  gives  a  mixture  of  pure  ice 
and  pure  salt  crystals.  Hence,  ice  and  salt  cannot  permanently 
exist  together  above  -21°.  Below  -6°  F.,  salt  will  no  longer  melt 
ice.  A  mixture  of  ice  and  salt,  giving  a  temperature  of  —  6°  F., 
is  called  a  freezing  mixture,  and  is  used  in  freezing  ice  cream  and 
ices. 

Molecular  weights  of  non-volatile  substances  can  be  measured  by 
simply  finding  out  what  weight  of  the  substance,  in  1000  c.c.  of 
water,  is  required  to  lower  the  freezing  point  from  0°  to  —1.86°. 

The    Vapor    Tension   of  Solutions:     Deliquescence.  —  A 

solute,  which  is  itself  nonvolatile,  tends  to  diminish  the  vapor 
tension  of  the  solvent.     It  hinders  the  emission  of  vapor. 
*  For  important  exceptions,  see  Chap.  XVI. 


SOLUTION  135 

If  the  substance  is  very  soluble,  and  the  solution  highly  concen- 
trated, the  lowering  in  the  vapor  tension  will  be  considerable.  In 
fact,  the  solution  may  give  a  vapor  pressure  of  water  less  than  that 
commonly  present  in  the  atmosphere.  Such  a  solution,  placed  in 
an  open  vessel,  will  not  evaporate.  On  the  contrary,  vapor  from 
the  air  will  enter  it,  and  it  will  increase  in  bulk.  For  this  reason, 
crystals  of  very  soluble  substances  are  usually  moist  and,  when 
exposed  to  the  air,  acquire  water  from  the  latter  and  dissolve  in 
this  water.  This  behavior  is  called  deliquescence,  and  is  exhibited, 
for  example,  by  the  hydrate  of  calcium  chloride  CaCl2,2H2O, 
which  is  consequently  used  for  drying  gases.  Magnesium  chloride 
MgCl2,  present  as  an  impurity  in  common  salt,  causes  the  latter  to 
become  moist  in  damp  weather. 

The  principle  involved  will  become  clear  if  we  imagine  two 
vessels,  one  containing  pure  water  and  one  an  aqueous  solution, 
to  be  placed  on  a  glass  plate  and  covered  by  a  bell  jar  (Fig.  61). 
Each  liquid  exchanges  water  molecules  with  the  moist  air  in  the 
jar,  but  the  solution  gives  off  water  more 
feebly  than  does  the  pure  water.  The 
result  is  that  the  latter  can  produce  a  pres- 
sure of  water  vapor  higher  than  that  which 
would  be  in  equilibrium  with  the  solution. 
The  solution,  therefore,  receives  continu- 
ously more  molecules  than  it  emits,  and 
increases  in  volume.  The  pure  water  thus  ^ 

gradually  passes  through  the  vapor  state 

into  the  solution  until  it  is  all  gone.  If  sufficient  water  was 
present,  the  process  would  go  on  until  the  solution  became  infinitely 
dilute. 

Boiling-Points  of  Solutions.  —  Since  the  solute  interferes 
with  the  emission  of  the  vapor  of  the  solvent,  it  naturally  makes 
the  solution  more  difficult  to  boil.  It  raises  the  boiling-point. 
Thus  one  gram-molecular  weight  of  sugar  (342  g.)  or  of  glycerine 
(92  g.),  dissolved  in  1000  c.c.  of  water,  will  elevate  the  boiling 
point  from  100°  to  100.52°  (for  exceptions,  see  Chap.  XVI). 

The  Laws  of  Osmotic  Pressure.  —  We  have  seen  that  a  dis- 
solved substance  exercises  a  pressure,  called  osmotic  pressure,  which 


136 


COLLEGE    CHEMISTRY 


is  proportional  to  the  concentration  of  the  solute  in  the  solvent. 
This  pressure  shows  itself  when  a  membrane  opposes  the  diffusion 
of  the  solute  from  the  solution  into  a  layer  of  pure  solvent.  The 
phenomena  can  be  illustrated  by  using  a  diffusion  shell,  of  test- 
tube  form,  attached  at  the  lower  end  of  a  glass  tube  (Fig.  62). 
We  may  place  the  solution  (e.g.,  of  sugar)  inside  the  shell,  and 
pure  water  outside.  The  material  of  the 
shell  is  such  that  the  water  can  pass  through 
it  in  either  direction,  but  the  molecules  of 
the  dissolved  substance  cannot  do  so.  Such 
a  membrane  is  called  semi-permeable. 

The  facts  observed  with  this  arrangement 
are  as  follows:  (1)  The  pure  solvent  passes 
into  the  solution.  Thus,  if  the  solution  is 
inside  the  shell,  water  enters  through  the  wall 
of  the  shell,  and  the  liquid  therefore  rises  in 
the  tube.  If  the  solution  is  outside,  water 
passes  out  of  the  shell,  and  the  liquid  in  the 
tube  falls.  (2)  If  solutions  of  different  con- 
centrations are  used  inside  and  outside,  the 
solvent  passes  from  the  more  dilute  into  the 
more  concentrated  solution,  so  that  the  tend- 
ency is  to  dilute  the  latter  and  concentrate 
the  former  until  both  have  the  same  con- 
centration. (3)  The  entrance  of  the  solvent 
can  be  prevented  by  the  application  of  pressure 
to  the  surface  of  the  liquid  in  the  tube.  With 
.  1  per  cent  sugar  solution  inside  and  water 

FlG-  62-  outside  (at  15°),  a  pressure  equal  to  500  mm. 

of  mercury  (0.66  atmos.)  per  square  centimeter  of  the  surface 
of  the  membrane  is  required  to  stop  the  entrance  of  the  sol- 
vent. This,  therefore,  is  the  value  of  the  so-called  osmotic  pres- 
sure. Since  the  entrance  of  the  solvent  is  due  to  the  dissolved 
substance,  and  the  solvent  is  really  drawn  forcefully  into  the 
solution,  it  might  be  more  appropriate  to  call  the  force  osmotic  suc- 
tion. Whatever  it  is  named,  however,  it  is  real,  and  its  value  can 
be  measured.  (4)  The  value  of  the  pressure  (or  suction)  increases 
in  proportion  to  the  absolute  temperature,  just  like  the  pressure 
of  a  gas.  (5)  The  value  of  the  pressure  (or  suction)  is  also  proper- 


SOLUTION  137 

tional  to  the  concentration  of  the  solute  in  the  solution.  Thus,  2  per 
cent  sugar  gave  (Pfeffer)  1016  mm.  pressure,  4  per  cent  sugar  2082 
mm.  (6)  When  different  solutes  are  compared,  it  is  found  that, 
at  the  same  temperature,  equal  numbers  of  molecules  of  the  solutes 
dissolved  in  equal  volumes  of  the  solvent  give  almost  equal  osmotic 
pressures  (or  suctions).  Thus,  one  mole  (342  g.)  of  sugar  C^H^On 
and  one  mole  (74  g.)  of  methyl  acetate  CH3C02CH3,  which,  in 
spite  of  the  great  difference  in  weight,  contain  equal  numbers  of 
molecules,  when  dissolved  separately,  each  in  ten  liters  of  water, 
give  in  both  cases  2.42  atmospheres  osmotic  pressure  (or  suction). 
(7)  Finally,  the  osmotic  pressure  (or  suction)  exercised  by  a  certain 
quantity  of  a  substance  in  solution  is  almost  identical  in  value  with 
the  gaseous  pressure  which  the  same  quantity  of  the  same  substance 
would  exercise  if  it  were  contained  as  a  gas  in  the  same  volume  at  the 
same  temperature.  Thus  44  g.  of  carbon  dioxide,  as  a  gas,  filling 
22.4  liters  (the  G.M.V.)  at  0°  exercises  one  atmosphere  gaseous 
pressure.  When  the  cube  is  filled  with  water,  and  the  gas  is 
thus  dissolved,  the  osmotic  pressure  of  the  solution  is  one 
atmosphere. 

These  facts  apply  to  substances  which  are  not  acids,  bases, 
or  salts.  We  shall  find  later  (see  Chap.  XVI)  that  the  osmotic 
pressures  of  the  members  of  these  three  classes  of  substances  are 
frequently  abnormally  high,  but  that  the  abnormality  is  easily 
explained. 

Osmotic  pressure  (or  suction)  is  a  subject  of  great  interest  in 
connection  with  the  physiology  of  plants  and  animals.  It  aids  in 
explaining  why  a  withered  flower,  containing  a  solution  in  its  cells, 
revives  when  placed  in  pure  water.  The  latter  enters  through  the 
walls  of  the  cells,  and  the  pressure  thus  produced  distends  the 
structure  and  stiffens  it.  Similarly,  the  ascent  of  the  water  from 
the  soil  into  the  roots  and  through  the  stem  of  a  growing  plant  is 
explained.  In  the  animal  body  also,  osmosis  plays  a  large  part. 

Measurements  of  osmotic  pressures  cannot  be  made  accurately 
with  a  diffusion  shell,  because  the  solute  is  able  to  some  extent  to 
pass  through  the  material.  Then,  too,  such  a  shell  can  be  used 
only  with  dilute  solutions,  because  it  will  not  withstand  high 
pressures.  In  making  accurate  measurements,  a  cell  of  porous 
porcelain  is  used,  and  the  pores  are  filled  with  a  gelatinous  precipi- 
tate of  cupric  ferrocyanide  Cu2Fe(CN)6  (q.v.). 


138  COLLEGE    CHEMISTRY 

Densities  of  Solutions.  —  The  density  or  specific  gravity  of 
a  solution  is  usually  greater  than  that  of  water  and,  in  each  case, 
varies  with  the  concentration.  For  commercial  purposes,  the 
concentrations  of  solutions  are  commonly  defined  by  the  specific 
gravity.  Thus,  we  purchase  ammonium  hydroxide  solution  of 
"0.88  sp.  gr.,"  meaning  35  per  cent  of  ammonia,  or  sulphuric  acid 
of  "1.84  sp.  gr.,"  meaning  94.8  per  cent  of  the  acid. 

The  commonly  greater  density  of  a  solution  is  utilized  in  making 
solutions  in  chemical  factories.  Shaking  several  tons  of  the 
mixture  is  out  of  the  question,  and  stirring  costs  money.  If  the 
solid  is  placed  in  the  bottom  of  the  tank,  under  water,  a  saturated 
solution  is  formed  in  the  lowest  layer  of  the  water,  and  passage  of 
the  dissolving  substance  into  the  upper  layers,  by  diffusion,  would 
take  months  or  years.  Hence  most  of  the  solid  would  remain  un- 
dissolved  (Fig.  57,  p.  127).  But  when  the  solid  is  placed  on  a 
shelf  near  the  surface  of  the  water,  the  solution  sinks  through  the 
water,  fresh  water  rises  to  the  shelf,  and  a  circulation  is  started. 
This  results  in  the  dissolving  of  the  whole  material  in  a  surprisingly 
short  time,  with  no  expenditure  of  labor  whatever. 

Changes  in  Volume  upon  Solution.  —  The  erratic  and,  at 
present,  unexplained  changes  in  volume  which  occur  when  a  sub- 
stance is  dissolved,  seem  to  indicate  that  the  process  is  less  simple 
than  we  have  thus  far  admitted,  and  that  chemical  changes  occur 
during  the  process..  Thus,  when  250  g.  of  common  salt  are  dis- 
solved in  1  liter  of  water  (=  1000  c.c.  =  1000  g.),  which  gives  a 
20  per  cent  solution,  the  volume  of  the  solution  is  only  1086  c.c. 
Since  the  250  g.  of  salt  occupied  116  c.c.  before  being  dissolved,  a 
shrinkage  of  1116  —  1086  or  30  c.c.  accompanied  the  process  of  solu- 
tion. On  the  other  han4,  214  g.  of  ammonium  chloride  (volume 
142.5  c.c.)  and  843.5  c.c.  of  water,  have  a  total  volume  of  986  c.c., 
but  when  dissolved  give  1000  c.c.  of  solution.  Here  there  is  an 
expansion  of  14  c.c.  Table  sugar,  however,  dissolves  in  water 
with  almost  no  change  in  volume. 

7s  Solution  a  Physical  or  a  Chemical  Change?  —  These 
phenomena  are,  in  part,  accounted  for  by  the  fact  that  water  is  not 
a  single  substance,  but  a  mixture.  It  is  largely  composed  of 
dihydrol  (H20)2,  with  much  trihydrol  (H20)3  near  to  0°  and  in- 
creasing quantities  of  monohydrol  H2O  at  higher  temperatures. 


SOLUTION  139 

When  any  substance  is  dissolved  in  considerable  amount  in  water, 
the  equilibrium  amongst  these  three  kinds  of  molecules  is  dis- 
turbed, and  their  proportions  change: 

2(H20)3  *=*  3(H20)2  *=*  6H20. 

Now,  equal  weights  of  these  three  kinds  of  water  occupy  different 
volumes,  and  hence  solution  is  accompanied  by  changes  in  the 
volume  of  the  water.  The  same  condition  in  water  explains  the 
point  of  maximum  density  (4°).  The  change  from  (H20)3  to 
(H20)2,  which  proceeds  as  the  temperature  rises  from  0°  to  4°,  is 
accompanied  by  a  shrinkage,  because  dihydrol  has  the  higher 
density.  Beyond  4°,  the  usual  expansion  with  rising  temperature 
prevails. 

There  is  also  evidence  that  many  dissolved  bodies  form  unstable 
compounds  with  water,  although  we  have  not  as  yet  definite  in- 
formation about  these  compounds. 

Dissolving  in  water  is,  therefore,  partly  a  chemical  and  only 
partly  a  physical  process  —  a  part  of  the  water  is  always  affected, 
and  a  part  or  all  of  the  solute  may  go  into  combination. 

Exercises.  —  1.   Give  other  examples  of  limited  solubility  in 

various  solvents  (p.  122). 

2.  What  weights  of  phosphoric  acid  (p.  94)  and  of  sodium 
hydroxide,  respectively,  are  required  to  make  1  liter  of  a  normal 
solution? 

3.  Express    the    concentrations    of    solutions    of    ammonium 
chloride,  saturated  at  0°  (sp.  gr.  1.076),  and  of  potassium  sulphate 
K2S04,  saturated  at  10°  (sp.  gr.  1.083),  in  terms  of  a  normal  solu- 
tion (p.  124). 

4.  Express  the  concentration  of  a  five  per  cent  aqueous  solution 
of  phosphoric  acid  (sp.  gr.  1.027),  in  terms  of  a  normal  and  a  molar 
solution,  respectively. 

5.  Explain  why,  (a)  pulverization  and,  (6)  agitation  hasten  the 
dissolving  of  a  solid  (cf.  pp.  331,  398). 

6.  Read  from  the  curves  (p.  131)  the  solubilities  of  potassium 
nitrate  at  15°,  of  potassium  chloride  at  30°,  of  potassium  chlorate 
at  45°.     What  are  the  relative  rates  at  which  the  solubilities  of 
these  salts  increase  with  rise  in  temperature? 

7.  If  5  g.  of  a  substance,  dissolved  in  1000  c.c.  of  water,  give  a 


140  COLLEGE    CHEMISTRY 

solution  freezing  at  —0.2°,  what  is  the  molecular  weight  of  the 
substance? 

8.  At  what  point  in  a  tank  of  water  should  you  introduce 
ammonia  gas,  in  order,  with  the  least  effort,  to  saturate  the  water? 
The  sp.  gr.  of  the  saturated  solution  is  0.88. 

9.  6  grams  of  a  substance  when  dissolved  in  200  c.c.  of  water 
give  a  boiling-point  of  102.6°.     What  is  the  molecular  weight  of 
the  substance? 

10.  1.6  grams  of  naphthalene  CioHg  when  dissolved  in  25  g.  of 
benzene  (freezing-point  5.48°)  gives  a  solution  which  freezes  at 
3.03°.     When  2.44  grams  of  another  substance  are  dissolved  in 
the  same  amount  of  benzene,  the  solution  freezes  at  3.52°.     What 
is  the  molecular  weight  of  the  latter  substance? 

11.  The  elevation  of  the  boiling-point  in  the  above  solution  of 
naphthalene  is  1.285°.     What  elevation  of  the  boiling-point  is  pro- 
duced in  the  second  solution? 


CHAPTER  XI 
HYDROGEN   CHLORIDE.     CALCULATIONS 

WE  have  had  occasion  several  times  to  mention  common  salt, 
or  sodium  chloride  NaCl.  This  is  one  of  the  most  familiar  chemi- 
cal substances.  Large  quantities  of  it  are  used  in  the  household, 
in  cooking  and  in  making  freezing  mixtures.  Still  larger  amounts 
are  consumed  in  manufacturing  washing  soda,  caustic  soda,  and 
soap,  for  all  of  which  it  furnishes  the  necessary  sodium.  It  is 
used  also  in  preserving  fish  and  other  foods.  It  supplies  the 
chlorine  used  in  bleaching  and  in  the  steriliza- 
tion of  city  waters.  We  shall  consider  it  first 
as  a  means  of  making  other  compounds  of 
chlorine. 

Preparation  of  Hydrogen  Chloride  HCl 
from  Salt.  —  When  some  concentrated  sul- 
phuric acid  is  poured  upon  sodium  chloride,  a 
vigorous  effervescence  is  noticed.  This  shows 
that  bubbles  of  a  gas  are  forming  upon  the 
salt  crystals  and  are  rising  through  the  acid 
and  bursting.  If  the  salt  be  placed  in  a 
flask  (Fig.  63),  the  acid  can  be  allowed  to  enter 
from  time  to  time  through  the  funnel.  When 
the  air  has  been  displaced  from  the  flask,  the  gas  issues  from 
the  delivery  tube.  If  the  correct  proportion  of  the  acid  is  used, 
and  only  a  gentle  heat  is  applied,  all  that  remains  in  the  flask  is 
a  white  solid,  sodium-hydrogen  sulphate  (sodium  bisulphate) 
NaHS04: 

NaCl  +  H2S04  ->  NaHS04  +  HCl  t  .*  (D 

*  The  arrow  directed  downwards  indicates  elimination  of  a  substance  by 
precipitation;  that  directed  upwards,  escape  as  a  gas  or  solution  of  a  solid. 

141 


142  COLLEGE    CHEMISTRY 

The  gas  is  extremely  soluble  in  water  and,  being  heavier  than 
air,  may  be  collected  by  upward  displacement  of  the  air  in  a  jar. 

The  action  described  is  the  one  which  occurs  in  the  laboratory. 
When  a  double  proportion  of  salt  and  a  high  temperature  are  used, 
a  second  action  occurs: 


NaCl  +  NaHSO4  ->  Na^  +  HC1  1 

and  sodium  sulphate  Na2S04  remains.  In  Europe  this  action  is 
employed,  with  furnace  heat,  in  manufacturing  sodium  sulphate, 
from  which  sodium  carbonate  is  afterwards  prepared.  The  hydro- 
gen chloride  passes  into  a  tower,  down  which  water  trickles  over 
lumps  of  coke,  and  is  dissolved.  The  aqueous  solution  is  called 
hydrochloric  acid  or,  in  commerce,  muriatic  acid  (Lat.,  brine 
acid). 

Hydrogen  Chloride  from  Other  Chlorides  'and  Other 
Acids.  —  The  chlorides  of  other  metals  could  be  substituted  for 
sodium  chloride  in  this  action,  and  all  the  more  soluble  ones  would 
give  hydrogen  chloride  freely.  Other  chlorides  are  all  more  ex- 
pensive, however,  than  is  common  salt. 

All  acids  contain  the  necessary  hydrogen  radical,  and  might 
offer  it  in  exchange  for  the  sodium  in  the  salt,  yet  in  practice  no 
other  acid  works  so  well  as  does  sulphuric  acid.  Concentrated 
phosphoric  acid  H3P04,Aq  acts  more  slowly,  giving  primary 
sodium  phosphate: 

NaCl  +  H3PO4  ->  NaH2P04  +  HC1  1  - 

The  Molecular  View  of  the  Interaction  of  Sulphuric  Acid 
and  Salt.  —  One  who  has  used  the  above-described  methods  for 
making  hydrogen  chloride  without  reflection  would  not  realize  the 
complexity  of  the  machinery  by  which  the  result  is  achieved.  The 
means  are  apparently  very  simple.  Yet  the  mechanical  features 
of  this  experiment,  when  laid  bare,  are  extremely  curious  and  in- 
teresting. A  single  fact  will  show  the  possibilities  which  are 
concealed  in  it. 

If  we  take  a  saturated  solution  of  sodium-hydrogen  sulphate  in 
water  and  add  to  it  a  concentrated  solution  of  hydrogen  chloride  in 
water  (concentrated  hydrochloric  acid),  we  shall  perceive  at  once 


HYDROGEN   CHLORIDE  143 

the  formation  of  a  copious  precipitate.  This  is  composed  entirely 
of  minute  cubes  of  sodium  chloride : 

NaHS04  +  HC1  ->  H2SO4  +  NaCl  J,  .*  (2) 

Now  this  action  is  nothing  less  than  the  precise  reverse  of  (1),  yet 
it  proceeds  with  equal  success.  In  fact,  this  chemical  interaction 
is  not  only  reversible  (pp.  93,  95),  but  can  be  carried  to  comple- 
tion in  either  direction.  It  is  only  in  presence  of  a  large  amount 
of  water  that  it  stops  midway  in  its  career  and  is  valueless  for 
securing  a  complete  transformation  in  either  direction: 

NaHS04  +  HC1  <=±  H2S04  +  NaCl. 

In  an  action  which  is  reversible,  if  the  products  remain  as  per- 
fectly mixed  and  accessible  to  each  other  as  were  the  initial  sub- 
stances, their  interaction  will  continually  undo  a  part  of  the  work 
of  the  forward  direction  of  the  change.  Hence,  in  such  a  case  the 
reaction  must,  and  does,  come  to  a  standstill  while  as  yet  only 
partly  accomplished;  but  this  was  not  the  case  with  actions 
(1)  and  (2).  Let  us  examine  the  means  by  which  the  premature 
cessation  of  each  was  avoided. 

In  equation  (1)  the  salt  dissolved  to  some  extent  in  the  sulphuric 
acid,  NaCl  (solid)  <=±  NaCl  (dslvd),  and  so,  by  contact  of  the  two 
kinds  of  molecules,  the  products  were  formed.  On  the  other  hand, 
the  hydrogen  chloride,  being  insoluble  in  sulphuric  acid,  escaped  as 
fast  as  it  was  formed:  HC1  (dslvd)  +±  HC1  (gas).  Hence,  in  that 
case,  almost  no  reverse  action  was  possible,  and  the  double  decom- 
position went  on  to  completion.  With  all  the  sodium-hydrogen 
sulphate  in  the  bottom  of  the  flask,  and  most  of  the  hydrogen 
chloride  in  the  space  above,  the  two  products  might  as  well  have 
been  in  separate  vessels  so  far  as  any  efficient  re-interaction  was 
concerned.  This  plan,  in  which  water  is  purposely  excluded,  forms 
therefore  the  method  of  making  hydrogen  chloride. 

In  equation  (2),  on  the  other  hand,  the  hydrogen  chloride  was 
taken  in  aqueous  solution,  and  was  mixed  with  a  strong  solution  of 
sodium  bisulphate.  The  acid  was,  therefore,  kept  permanently  in 
full  contact  with  the  sodium  bisulphate.  It  had  in  this  case,  every 
opportunity  to  interact  with  the  latter  and  no  chance  of  escape. 
Every  molecule  of  each  ingredient  could  reach  every  molecule  of 
*  See  footnote  to  p.  141. 


144  COLLEGE    CHEMISTRY 

the  other  with  equal  ease.  Furthermore,  the  sodium  chloride, 
produced  as  a  result  of  their  activity,  is  not  very  soluble  in  con- 
centrated hydrochloric  acid  (far  less  so  than  in  water),  and  so  it 
came  out  as  a  precipitate:  NaCl  (dslvd)  <=±  NaCl  (solid).  But 
this  was  almost  the  same  as  if  it  had  gone  off  as  a  gas.  It  meant 
that  the  greater  part  of  the  salt  was  in  the  solid  form.  It  was  in 
a  state  of  fine  powder,  it  is  true.  But,  in  the  molecular  point  of 
view,  the  smallest  particle  of  a  powder  contains  millions  of  mole- 
cules, and  most  of  these  are  necessarily  buried  in  the  interior  of  a 
particle.  Thus,  the  sodium  chloride  was  no  longer  able  to  interact 
effectively  molecule  to  molecule  with  the  other  product,  the  sul- 
phuric acid.  Hence,  there  was  little  reverse  action  to  impede  the 
progress  of  the  primary  one.  Thus  (2)  is  nearly  as  perfect  a  way 
of  liberating  sulphuric  acid  as  (1)  is  of  liberating  hydrogen 
chloride. 

This  discussion  is  given  to  illustrate  the  displacement  of  a  chemi- 
cal equilibrium,  and  to  explain  the  method  of  preparing  hydrogen 
chloride.  It  also  throws  an  interesting  light  on  chemical  affinity, 
however.  Considering  action  (1),  by  itself,  we  might  reason  that 
the  hydrogen  chloride  was  formed  because  the  affinity  of  the  hydro- 
gen (H)  for  chlorine  (Cl)  was  greater  than  for  the  sulphate  radical 
(804).  But,  if  we  did  so,  then  in  action  (2)  we  should  be  compelled 
to  reason  similarly  that  the  preponderance  of  affinity  was  just  the 
opposite.  In  point  of  fact,  no  conclusion  about  relative  affinity 
can  be  drawn  from  these  actions.  The  effects  of  affinity  are  here 
entirely  subordinated  by  the  effects  of  a  purely  mechanical  ar- 
rangement, depending  on  solubility.  When  the  activities  of  the 
acids  are  properly  compared,  hydrochloric  acid  is  found  to  be 
considerably  more  active  than  sulphuric  acid. 

Physical  Properties.  —  Hydrogen  chloride  is  a  colorless  gas, 
which  produces  a  suffocating  effect  when  inhaled. 

Density  (H  =  1),  18.23.  Grit,  temp.,  +52°. 

Weight  of  22.4  1.,  36.73  g.  Boiling-point  (liq.),  -83.7°. 

Sol'ty  in  Aq  (0°),  50,300  vols.  in  100.  Melting-point  (solid),  -110°. 

The  gas  is  one-fourth  heavier  than  air.  On  account  of  its  great 
solubility,  when  it  streams  into  the  air  it  condenses  atmospheric 
moisture  into  a  fog  (of  drops  of  hydrochloric  acid).  The  extreme 


HYDROGEN   CHLORIDE  145 

solubility  may  be  shown  by  filling  a  dry  flask  (Fig.  64)  with  the 
gas.  The  " dropper"  contains  water,  and  is  closed  at  the  tip 
with  soft  wax.  A  drop  of  water,  expelled  by  pinching  the  "drop- 
per," dissolves  so  much  of  the  gas  that  the  water  is  forced  in  by 
atmospheric  pressure,  like  a  fountain,  through  the 
longer  tube. 

Both  in  the  gaseous  and  liquefied  states  it  is  a  non- 
conductor of  electricity.  Its  heat  of  solution  is 
17,400  calories  (p.  85).  On  account  of  its  high 
concentration,  the  saturated,  aqueous  solution  may 
be  looked  upon  as  a  mixture  of  liquefied  hydrogen 
chloride  and  water. 

When  the  concentrated  aqueous  solution  is  heated,  it 
is  the  gas  and  not  the  water  which  is  driven  out,  for 
the  most  part.  When  the  concentration  has  been 
reduced  to  20.2  per  cent,  the  rest  of  the  mixture 
distils  unchanged  at  110°.  This  occurs  because,  at 


this  concentration,  the  gas  is  carried  off  in  the  bubbles  FlG-  64. 
of  steam  in  the  same  proportion  in  which  it  is  present  in  the 
liquid.  If  a  dilute  solution  is  used,  water  is  the  chief  product  of 
distillation  (about  100°),  but  gradually  the  boiling-point  rises  and, 
when  the  concentration  has  reached  20.2  per  cent  once  more, 
the  same  hydrochloric  acid  of  constant  boiling-point  (110°  at  760 
mm.),  as  it  is  called,  forms  the  residue. 

Chemical  Properties.  —  Hydrogen  chloride  is  extremely  stable, 
as  we  might  expect  from  the  vigor  with  which  the  elements  of  which 
it  is  composed  combine  (see  p.  160).  On  being  heated  to  a  tempera- 
ture of  1800°,  however,  it  begins  to  dissociate  into  its  constituents. 

In  the  chemical  point  of  view,  it  is  on  the  whole  rather  an  indif- 
ferent substance.  Hydrogen  chloride  (the  gas)  has  no  action  upon 
any  of  the  non-metals,  such  as  phosphorus,  carbon,  sulphur,  etc. 
Many  of  the  metals,  however,  particularly  the  more  active  ones, 
such  as  potassium,  sodium,  and  magnesium,  decompose  it.  Hy- 
drogen is  set  free,  and  the  chloride  of  the  metal  is  formed.  The 
equation  representing  the  weights,  is  K  +  HC1  — »  KC1  +  H.  But 
the  molecular  formula  (p.  Ill)  of  hydrogen  is  H2,  hence  the  cor- 
rect equation  is: 


146  COLLEGE    CHEMISTRY 

Hydrogen  chloride  unites  directly  with  ammonia  gas  to  form  a 
cloud  of  solid  particles  of  ammonium  chloride  (HC1  -f-  NH3  —  > 


Chemical  Properties  of  Hydrochloric  Acid.  —  The  solution 
of  hydrogen  chloride  in  water  is  an  entirely  different  substance  in 
its  chemical  behavior  from  hydrogen  chloride.  It  is  strongly  acid, 
turning  blue  litmus  red.  The  gas  and  liquefied  gas  have  no  such 
property.  The  solution  conducts  electricity  very  well,  and  is  de- 
composed in  the  process  (p.  55),  giving  hydrogen  at  the  negative 
wire  and  chlorine  at  the  positive  wire: 

2HC1-+H2  (neg.  wire)  +  C12  (pos.  wire). 

The  gas  and  the  liquefied  gas  are  practically  nonconductors. 

The  metals  preceding  hydrogen  in  the  order  of  activity  (p.  60), 
when  introduced  into  hydrochloric  acid,  displace  the  hydrogen 
(p.  55),  and  form  the  chloride  of  the  metal.  In  the  case  of  zinc 
the  action  was  represented  by  the  equation: 

Zn  +  2HCl-+ZnCl2  +  H2. 

The  aqueous  solution  of  hydrogen  chloride  interacts  rapidly  with 
most  oxides  and  hydroxides  of  metals,  as,  for  example,  those  of  zinc  : 

ZnO  +  2HC1  -»  ZnCl2  +  H2O, 
Zn(OH)2  +  2HC1  ->  ZnCl2  +  2H2O. 

Here  no  free  hydrogen  is  obtained,  since  the  oxygen  in  the  oxide, 
and  the  hydroxyl  in  the  hydroxide,  unite  with  it  to  form  water. 
In  each  case,  however,  the  chloride  of  the  metal  is  obtained.  It 
may  be  noted,  in  passing,  that  all  acids  behave  in  a  similar  manner 
towards  oxides  and  hydroxides  of  metals,  giving  water  and  a  com- 
pound corresponding  to  the  chloride.  Dilute  sulphuric  acid,  for 
example,  gives  sulphates. 

Modes  of  Preparing  Chlorides.  —  In  the  preceding  section 
three  kinds  of  actions,  each  constituting  a  different  mode  of  pre- 
paring chlorides,  have  been  mentioned  incidentally.  There  are 
two  others.  The  simplest  is  the  direct  union  of  the  element  with 
chlorine  (Zn  +  C12  -*  ZnCl2)  .  The  other  method  is  illustrated 
in  the  case  of  the  precipitation  of  silver  chloride  by  adding  a  solu- 


HYDROGEN   CHLORIDE  147 

tion  of  a  chloride  to  a  solution  of  silver  nitrate.  Here  the  forma- 
tion of  the  chloride  occurs  by  exchange  of  another  radical  (p.  53) 
for  the  chloride  radical: 

AgN03  +  NaCl  ->  AgCl  [  +  NaN03. 

The  insoluble  chlorides  (see  p.  164)  can  be  made  conveniently  by 
this  plan.  The  formation  of  the  precipitates,  for  example  that  of 
silver  chloride,  is  used  as  a  test  for  the  presence  of  a  soluble  chlo- 
ride in  the  solution. 

Uses  of  Hydrochloric  Acid.  —  This  substance  is  used,  in 
Europe,  as  a  commercial  source  of  chlorine.  It  is  employed  in 
cleaning  metals,  and  in  the  manufacture  of  chlorides  of  metals. 
It  is  an  important  component  of  the  gastric  juice  of  the  stomach, 
although  the  proportion  is  only  about  1  part  in  500. 

Precipitation.  —  When  two  soluble  substances  are  dissolved, 
separately,  and  the  solutions  are  mixed,  chemical  interaction  fre- 
quently occurs,  as  in  the  case  of  salt  and  silver  nitrate  (see  also 
p.  143).  If  one  of  the  products  is  insoluble,  then  a  supersaturated 
solution  of  this  product  is  at  once  produced.  As  a  rule,  this  sub- 
stance almost  immediately  becomes  visible  as  a  fine  powder,  called 
a  precipitate,  suspended  in  the  liquid. 

The  insoluble  product  can  often  be  recognized  by  its  physical 
appearance,  and  so  this  sort  of  action  is  frequently  used  as  a 
test  for  one  of  the  original  substances.  Thus  many  precipitates 
have  a  distinctive  color.  Again,  precipitates  which  are  colorless, 
or  have  the  same  color,  differ  in  appearance,  and  are  described  as 
gelatinous,  curdy,  pulverulent,  or  crystalline.  In  the  first  two  cases, 
the  precipitation  is  so  sudden  that  there  is  no  time  for  crystals  to 
be  formed,  and  the  product  is  amorphous  (Gk.,  without  form). 
Thus  silver  chloride  is  curdy,  and  precipitated  sodium  chloride 
(p.  143)  is  crystalline. 

Fourth  Variety  of  Chemical  Change:  Double  Decompo- 
sition.—  In  this  chapter  we  encounter  for  the  first  time  the 
fourth  variety  of  chemical  change.  Upon  examining  the  equa- 
tion for  the  action  of  sodium  chloride  and  silver  nitrate,  we  see 
that  the  silver  nitrate  decomposed  into  its  radicals  (Ag)  and  (NO3). 


148  COLLEGE    CHEMISTRY 

The  sodium  chloride  also  decomposed  into  its  radicals  (Na)  and 
(Cl).  The  (Ag)  then  united  with  the  (Cl)  and  the  (Na)  with  the 
(N03). 

AgN03  +  NaCl  ->  AgCl  +  NaN03. 

Since  both  of  the  original  substances  decomposed,  this  is  called  a 
double  decomposition.  An  exchange  of  radicals  occurred. 

The  action  by  which  hydrogen  chloride  was  prepared  (p.  142) 
belonged  to  the  same  class : 

NaCl  +  HHSO4  ->  NaHSO4  +  HC1. 

Double  decompositions  involving  acids,  bases,  and  salts  are  all 
reversible  reactions.  The  fact  that  many  of  them  proceed,  never- 
theless, to  practical  completion  has  already  been  explained  at  length 
(pp.  142-144). 

The  Varieties  of  Chemical  Change.  —  Most  chemical 
changes  belong  to  one  of  the  four  varieties : 

1.  Combination,  e.g.,  Fe  +  S  — >  FeS. 

2.  Decomposition,  e.g.,  2KC103  —>  2KC1  +  302. 

3.  Displacement,  e.g.,  Zn  +  2HC1  ->  H2  +  ZnCl2. 

4.  Double  Decomposition,  e.g.,  AgN03  +  HC1  -» AgCl  +  HN03. 
In  the  first,  2  (or  more)  substances  give  1  substance. 

In  the  second,  1  substance  gives  2  (or  more)  substances. 

In  the  third,  1  element  and  1  compound  give  1  element  and  1 
compound. 

In  the  fourth,  2  compounds  give  2  compounds. 

Occasionally,  one  compound  gives  one  (different)  compound,  a 
change  called  internal  rearrangement.  Nearly  all  chemical  changes, 
so  far  as  their  mechanism  is  concerned,  can  be  classified  under 
one  or  other  of  these  five  kinds. 

A  dissociation  (p.  93)  is  both  a  combination  and  a  decomposi- 
tion, because  it  is  reversible.  For  example: 

2H2O  <±  2H2  +  O2. 

Electrolysis  is  decomposition  by  an  electric  current. 

The  foregoing  varieties  of  chemical  action  are  general,  and  not 
limited  to  any  classes  of  elements.  Oxidation  (p.  36)  and  reduc- 
tion (p.  37)  are  so  limited.  Thus  combination  with  oxygen  is 
oxidation,  while  combination  with  hydrogen  is  reduction.  A  more 


CALCULATIONS  149 

complete  discussion  of  action  of  these  classes  will  be  given  later 
(see  Chapter  XXIII). 

The  reader  should  classify  each  action  mentioned  in  the  text, 
and  so  become  familiar  with  the  chemical  point  of  view  which  this 
classification  represents. 

Salts.  —  We  have  seen  that  an  acid  contains  hydrogen  H  as 
a  radical  (p.  52),  and  a  base  contains  the  radical  hydroxyl  OH 
(p.  94).  The  name  salts  is  given  to  the  class  of  substances  which 
contain  a  positive  and  a  negative  radical,  neither  of  which  is  hydro- 
gen nor  hydroxyl.  For  example,  NaCl,  Na2S04,  AgNOs  are  the 
formulae  of  salts.  Salts  are  so  named  because  they  resemble 
common  salt  in  having  two  radicals,  and  entering  readily  into 
double  decomposition. 

Sodium-hydrogen  sulphate  NaHSCX  is  classed  as  an  acid  salt, 
because  it  has  a  positive  and  a  negative  radical,  and  a  hydrogen 
radical  in  addition. 

CALCULATIONS 

Familiarity  with  the  interpretation  of  molecular  equations  is 
best  obtained  by  making  simple  calculations  based  upon  their 
common  uses  in  chemistry. 

Weights.  —  When  a  problem  in  regard  to  weights  of  material 
used  or  produced  in  a  given  action  is  to  be  solved,  the  molecular 
equation  is  to  be  written  and  the  weights  inserted  beneath  the 
formulae.  The  mode  of  calculation  has  been  described  already 
(pp.  67,  116). 

Weights  and  Volumes.  —  When  a  problem  involving  weights 
and  volumes  is  to  be  solved,  the  molecular  equation  is  to  be  written, 
and  both  the  weights  and  volumes  are  to  be  inserted.  Note,  how- 
ever, that  only  the  volumes  of  the  substances  in  the  gaseous  condi- 
tion are  considered. 

For  example,  what  volume  of  oxygen  is  obtained  from  60  g.  of 
potassium  chlorate?  The  molecular  equation,  made  as  already 
described  (p.  116),  together  with  the  full  interpretation,  are  as 
follows : 


150  COLLEGE    CHEMISTRY 

2KC103  ->        2KC1  +     302. 

f2  (39.1  +  35.46  +  48)          2  (39.1  +  35.46)  3  X  32 

WEIGHTS:  |  -  _T_  -  H9.1  g.  ^96^ 

VOLUMES:  3  X  22.4  1.     , 

Observe  that  no  volumes  are  given  under  the  chlorate  and  chlo- 
ride of  potassium.  This  is  because  their  volumes  in  the  gaseous 
condition  can  be  of  no  practical  use,  since  they  are  solids  which  are 
melted,  but  not  vaporized  during  this,  or  any  action  in  which  we 
employ  them.  Now,  as  to  the  problem  in  hand,  it  is  concerned 
with  a  weight  of  potassium  chlorate  and  a  volume  of  oxygen. 
Reading  from  the  equation,  our  information  on  these  points  is 
that  245.1  g.  of  potassium  chlorate  give  67.2  liters  (observe  that 
the  coefficients  are  used,  as  well  as  the  molecular  weights,  in  these 
numbers)  of  oxygen  at  0°  and  760  mm.,  and  the  question  is:  What 
volume  will  60  g.  give?  By  proportion,  245.1  g.  :  67.2  1.  :  :  60  g.  :  x  1., 
where  x  =  16.45  liters.  If  a  different  temperature  and  pressure 
had  been  specified,  either  the  volume  in  the  equation,  or  the  an- 
swer, would  have  had  to  be  converted,  by  rule,  to  the  given  condi- 
tions. 

It  saves  time  not  to  write  out,  as  above,  the  whole  interpreta- 
tion, but  only  the  parts  required.  For  example,  if  the  question 
is:  What  volume  of  chlorine  is  needed  to  give  25  g.  of  aluminium 
chloride?  we  may,  if  we  choose,  omit  all  the  data  excepting  the 
volume  of  the  chlorine  and  the  weight  of  the  aluminium  chloride, 
thus: 

2A1  +      3C12       ->       2A1C13 
3  X  22.4  1.  2  X  133.5  g. 

The  volume  of  chlorine  required  is  25  X  3  X  22.4  -f-  (2  X  133.5) 
liters.  These  illustrations  show  the  method  of  calculating  actual 
volumes  (see  Exercises  1,  2). 

Relative  Volumes  Alone.  —  If  the  question  concerns  relative 
volumes  only,  then  it  is  simplest  to  use  the  interpretation  of  the 
equation  in  terms  of  molecules.  For  example:  What  relative 
volumes  of  hydrogen  chloride  and  oxygen  are  required  in  Deacon's 
process  (see  p.  155)?  The  molecular  equation  is 


02->2H2O  +  2C12. 
MOLECULES:  4  1  2 


CALCULATIONS  151 

Since  equal  numbers  of  molecules  of  gases  occupy  equal  volumes, 
the  proportion  4  molecules  of  hydrogen  chloride  to  1  molecule  of 
oxygen  shows  the  ratio  to  be  4  :  1  by  volume.  Similarly,  every  4 
molecules  of  hydrogen  chloride  give  2  molecules  of  chlorine,  so  that 
the  ratio  of  these  substances  by  volume  is  4:  2,  or  2  :  1. 

In  regard  to  the  water,  since  that  is  not  a  gas  at  common  tem- 
peratures, the  question,  if  asked,  must  be  more  specific  :  What  are 
the  relative  volumes  of  steam  and  chlorine  in  the  product,  as  com- 
monly delivered  by  this  action  at  400°?  It  is  2:2,  or  1:1. 
What  are  the  relative  volumes  of  water  and  chlorine,  after  the 
products  have  cooled  to  room  temperature?  The  water  is  no 
longer  a  gas,  so  that  it  occupies,  relatively,  almost  no  volume.* 

What  is  the  total  volume-change  in  the  foregoing  action  above 
100°?  It  is  a  change  from  5  molecules  to  4.  The  volume  changes 
in  the  same  ratio.  But  at  0°  the  volume-change  is  from  5  volumes 
to  2,  for  the  water  does  not  appreciably  add  to  the  volume  of  the 
products  (see  Exercises  3,  4). 

Relative  Volumes,  Again.  —  When  we  know  the  molecular 
formulae  of  the  single  substances  concerned  in  an  action,  the  equa- 
tion can  be  made,  and  the  relative  volumes  determined,  without 
actual  measurement.  For  example  :  What  volume-change  will  be 
observed  when  a  mixture  of  carbon  monoxide  and  oxygen  has  ex- 
ploded, and  the  temperature  has  once  more  reached  that  of  the 
room?  The  molecular  formulae  are  CO,  02,  and  C02.  The  equa- 
tion representing  the  weights  is  CO  +  0  —  >  CO2.  The  molecule 
of  oxygen,  however,  being  O2,  we  cannot  employ  less  than  this 
quantity  in  a  molecular  equation,  so  that  the  equation  becomes: 

02->2C02. 


Three  molecules,  therefore,  give  two,  throughout  the  whole  mass, 
and  therefore  three  volumes  will  become  two,  if  the  pressure  and 
temperature  are  the  same  at  the  beginning  and  end  of  the  action. 
*  Of  course  if  an  exact  answer  must  be  given,  it  can  be  given.  But  for 
this  we  require  the  weight  and  specific  gravity  of  the  product.  Thus,  2H2O 
represents  2  X  18  g.  of  water.  The  sp.  gr.  of  water  is  1.  Therefore  the 
volume  of  water  formed  is  36  c.c.  The  volume  of  2C12  is  2  X  22.4,  or  44.8 
liters  at  0°.  The  ratio  of  water  to  chlorine  by  volume  at  0°  is  therefore  36: 
44,800.  But,  as  a  rule,  we  simply  give  the  volumes  of  solids  and  liquids  as 
zero,  compared  with  those  of  the  gases  concerned  in  the  same  action. 


152  COLLEGE    CHEMISTRY 

If  we  remember  that  all  volatile  compounds  of  carbon  and 
hydrogen  burn  to  form  water  and  carbon  dioxide,  the  molecular 
equation  for  any  such  combustion  may  easily  be  made,  and  the 
volumes  of  all  the  materials  ascertained.  When  water  is  a  product, 
only  its  volume  as  steam  is  given  by  the  equation  (see  Exercises 
4,5). 

Relative  Densities  of  Gases.  —  Knowing  by  heart  the  molec- 
ular formulae  of  gaseous  substances,  as  we  must  know  them  for 
many  purposes,  it  is  unnecessary  to  burden  our  minds  with  other 
data  in  regard  to  the  relative  weights  of  gases.  Is  hydrogen  chloride 
(HC1)  heavier  or  lighter  than  carbon  dioxide  (CO2)?  These  for- 
mulae represent  the  weights  of  equal  volumes  (22.4 1.),  namely,  36.46 
g.  and  44  g.,  respectively.  Hence  the  former  gas  is  a  little  lighter. 

Remembering  that  the  G.M.V.  of  air  weighs  28.955  g.  (Table, 
p.  101),  we  can  compare  the  weight  of  any  gas  with  that  of  air  in 
the  same  way.  What  are  the  relative  weights  of  acetylene  (C2H2, 
p.  105)  and  sulphur  dioxide  (S02)  as  compared  with  air?  The 
G.M.V.  cube  holds  formula- weights  of  the  first  two,  namely  26  g. 
and  64  g.,  and  28.955  g.  of  air.  Hence  acetylene  is  a  little  lighter 
than  air,  and  sulphur  dioxide  more  than  twice  as  heavy  (see 
Exercise  6). 

Exercises.  —  1.  What  volume  of  oxygen  at  10°  and  750  mm.  is 
obtainable  by  heating  50  g.  of  potassium  chlorate  (pp.  116,  150)? 

2.  What  volume  of  oxygen  at  20°  and  760  mm.  is  required  to 
convert  16  g.  of  iron  into  dehydrated  rust  (Fe2Os)  (p.  150)? 

3.  Write  out  the  molecular  equations  for  the  interactions  of 
methane  and  chlorine  giving  CH3C1;    and  for  the  burning  of 
phosphorus  (vapor)  in  oxygen  (p.  105).     Deduce  the  volume  re- 
lations of  the  initial  substances,  and  of  the  products,  at  various 
temperatures  in  each  case. 

4.  Write  out  the  molecular  equations  for  the  interactions  of 
acetylene  and  oxygen  (p.  105),  and  of  alcohol  vapor  (b.-p.  78°)  and 
oxygen.     Deduce  the  volume  relations  of  the  initial  substances  and 
of  the  products  at  0°  and  at  100°  in  each  case. 

f      5.   The  molecular  weight  of  cyanogen  is  52.08.     What  is  its  den- 

\    sity  referred  to  air,  'and  what  the  weight  of  1 1.  at  0°  and  760  mm.? 

It  contains  46.08  per  cent  carbon  and  53.92  per  cent  nitrogen. 


CALCULATIONS 


at  is  the  formula  of  the  substance  (p.  45)?  Exploded  with 
oxygen  it  forms  carbon  dioxide  and  free  nitrogen.  What  will  be 
the  relative  volumes  of  the  materials  before  and  after  the  inter- 
action (p.  151?) 

6.  What  are  the  relative  weights  of  equal  volumes  of  hydrogen 
sulphide  (H2S),  and  hydrogen  iodide  (HI),  compared  with 


CHAPTER  XII 
CHLORINE 

CHLORINE  was  first  recognized  as  a  distinct  substance  by  Scheelc 
(1774).  He  obtained  it  from  salt  by  means  of  manganese  dioxide, 
using  the  method  described  below.  It  was  supposed  to  be  a  com- 
pound containing  oxygen  until  Davy  (1809-1818)  demonstrated 
that  it  was  an  element. 

Occurrence.  —  Chlorine  does  not  occur  free  in  nature.  There 
are,  however,  many  compounds  of  it  to  be  found  in  the  mineral 
kingdom.  Sea-water  contains  a  number  of  chlorides  in  solution. 
Of  the  3.6  per  cent  of  solid  matter  in  sea-water,  nearly  2.8%  is 
sodium  chloride  NaCl.  During  past  geological  ages  the  evapora- 
tion of  sea-water  has  led  to  the  formation  of  immense  deposits  of 
the  compounds  usually  found  in  such  water.  Thus,  at  Stassfurt, 
such  strata  attain  a  thickness  of  over  a  thousand  feet.  Certain 
layers  of  these  strata  are  composed  mainly  of  sodium  chloride 
(rock  salt).  In  other  layers  potassium  chloride  (sylvite),  an  in- 
dispensable fertilizer,  and  other  compounds  of  chlorine,  occur. 

Preparation.  —  Chlorine  cannot  be  obtained  with  the  same 
ease  as  oxygen.  There  are  only  a  few  chlorides,  such  as  those  of 
gold  and  platinum,  which  lose  chlorine  when  heated,  and  they  are 
too  expensive  or  difficult  to  make  for  laboratory  use.  We  employ 
therefore  methods  like  those  used  for  the  preparation  of  hydrogen 
(cf.  p.  53).  We  may  (1)  decompose  any  chloride  by  means  of 
electricity,  just  as,  to  get  hydrogen,  we  electrolyzed  a  dilute  acid 
(p.  55).  Or  (2)  we  may  take  some  inexpensive  compound  of 
chlorine,  such  as  hydrogen  chloride  (HC1),  and  by  means  of  some 
simple  substance  which  is  capable  of  uniting  with  the  other  con- 
stituent —  here  oxygen  serves  the  purpose  —  secure  the  liberation 
of  the  element.  Or  (3)  —  and  this  turns  out  to  be  the  most  con- 
venient laboratory  method  —  we  may  use  a  more  complex  action. 

154 


CHLORINE 


155 


Electrolysis  of  Chlorides.  —  Hydrogen  chloride  and  those 
chlorides  of  metals  which  are  soluble  in  water  are  all  decomposed 
when  a  current  of  electricity  is  passed  through  the  aqueous  solu- 
tion. They  yield  chlorine  at  the  positive  electrode.  The  other 
constituent,  the  hydrogen  (Fig.  65),  manganese,  or  whatever  it  may 
be,  is  liberated  at  the  negative 
wire.  Since  the  chlorine  is  solu- 
ble in  water,  the  effervescence 
due  to  its  release  is  not  notice- 
able until  the  liquid  round  the 
electrode  has  become  saturated 
with  the  gas:  C12  (dslvd)  <=»  C12 
(gas).  The  shape  of  the  appa- 
ratus keeps  the  two  products 
from  mingling.  The  presence 
of  the  chlorine  in  the  liquid  at 
the  positive  end  may  be  shown 
by  a  suitable  test  (p.  161). 

In  commerce  chlorine  is  now 
obtained  chiefly  by  this  method, 
sodium  chloride  or  potassium 
chloride  being  the  source  of  the  element.  Electrodes  of  artificial 
graphite  are  used,  as  most  other  conductors  unite  with  the 
chlorine.  The  potassium  or  sodium,  as  the  case  may  be, 
travels  towards  the  negative  electrode,  but  is  not  liberated. 
Instead,  potassium  or  sodium  hydroxide  (q.v.)  accumulates  in 
the  solution  round  the  plate  and  hydrogen  escapes.  The  chlo- 
rine is  released  at  the  positive  electrode,  as  usual.  The  hydro- 
gen, the  hydroxide  and  the  chlorine  all  find  commercial  applica- 
tions. The  chlorine  is  either  liquefied  by  compression  in  steel 
cylinders  or  is  employed  at  once  for  making  bleaching  powder 
(see  index). 

Action  of  Free  Oxygen  on  Chlorides.  —  Sodium  chloride  is 
the  cheapest  source  of  chlorine,'but  oxygen  does  not  interact  with 
it  even  at  a  high  temperature.  By  treating  the  sodium  chloride 
with  sulphuric  acid,  therefore,  the  chlorine  is  first  transferred  into 
combination  with  the  hydrogen  of  the  acid,  giving  hydrogen 
chloride  (p.  141).  In  order  to  liberate  chlorine  from  the  hydrogen 


FIG.  65. 


156  COLLEGE    CHEMISTRY 

chloride,  we  may  then  combine  the  hydrogen  with  oxygen  obtained 
from  the  air. 

Skeleton:  HC1  +  O  ^±  H2O  +  Cl. 

Balanced:  2HC1  +  0  *±  H20  +  2C1. 

Molecular:  4HC1  +  O2  ?±  2H2O  +  2C12. 

The  two  gases  interact  so  slowly,  however,  that  a  contact  agent 
must  be  employed.  The  mixture  of  air  and  hydrogen  chloride  is 
passed  over  pieces  of  heated  pumice-stone  (Fig.  66)  or  broken 
brick  previously  saturated  with  cupric  chloride  solution.  A  tem- 
perature of  about  370°  is  used. 
Furthermore,  the  action  is  re- 
versible (read  the  equation 
backwards)  and  equilibrium  is 

reached  when  80  per  cent  of  the  hydrogen  chloride  has  been 
decomposed.  Hence  20  per  cent  of  this  gas  passes  on  unchanged. 
Only  80  per  cent  of  the  hydrogen  chloride  and  oxygen  are  changed 
into  steam  and  chlorine,  because  the  latter  substances  are  continu- 
ously interacting  to  reproduce  hydrogen  chloride  and  oxygen.  If 
one  substance  could  be  separated  (p.  143)  from  the  other,  to  pre- 
vent the  backward  action,  the  yield  would  be  raised  to  100  per 
cent.  In  the  product,  the  chlorine  is  mixed  with  steam  and  with  a 
very  large  volume  of  nitrogen  which  entered  with  the  oxygen,  as 
well  as  with  unused  hydrogen  chloride,  so  that,  for  making  the 
pure  substance,  this  method  (Deacon's  process)  is  quite  unsuitable. 
Bleaching  powder,  however,  can  be  made  by  its  means. 

The  relative  volumes  in  this  reaction  (see  p.  150)  are  indicated 
by  the  numbers  of  molecules  in  the  equation.  Four  volumes  of 
hydrogen  chloride  and  one  volume  of  oxygen  give  two  volumes  of 
steam  and  two  volumes  of  chlorine. 

The  above  action  is  spoken  of  as  an  oxidation.  It  is  true  that 
no  oxygen  is  actually  introduced  into  the  hydrogen  chloride  as 
a  whole.  The  removal  of  hydrogen  from  combination  with  the 
chlorine  is,  however,  the  first  step  towards  the  introduction  of 
oxygen  into  combination  with  the  latter,  and  is  essentially  an 
oxidation. 

Action  of  Combined  Oxygen  upon  Chlorides.  —  The  best 
laboratory  method  for  making  chlorine  is  to  place  some  solid 


CHLORINE 


157 


FIG.  67. 


potassium  permanganate  in  a  flask,  arranged  like  that  in  Fig.  67. 
Concentrated  hydrochloric  acid  (an  aqueous  solution  of  hydrogen 
chloride),  diluted  with  one- 
third  of  its  volume  of  water, 
is  allowed  to  fall  upon  the  com- 
pound drop  by  drop  from  the 
dropping  funnel.  The  action 
is  very  rapid,  the  acid  is  ex- 
hausted almost  as  fast  as  it 
falls,  and  so  the  stream  of  gas 
can  be  stopped  by  simply  clos- 
ing the  stopcock.  The  gas 
is  passed  through  a  washing 
bottle  containing  water,  in. 
order  to  remove  any  hydrogen 
chloride  which  may  be  carried 
over.  It  may  be  dried,  if 
necessary,  in  a  second  washing  bottle  containing  concentrated 
sulphuric  acid.  It  cannot  be  collected  over  water  on  account  of 
its  solubility,  so  that  jars  are  usually  filled  with  it  by  upward 
displacement  of  air. 

Skeleton:     KMn04  +  HC1  ->  H20  +  KC1  +  MnCl2  -f  Cl. 

The  O4,  being  all  converted  into  water,  requires  8H,  and  therefore 
8HC1,  for  the  action.  The  two  metals,  potassium  and  manganese, 
give  their  respective  chlorides,  KC1  and  MnCl2.  This  uses  3C1,  and 
hence  5C1  remains  over  to  be  liberated: 

Balanced:       KMnO4  +  8HC1  ->  4H20  +  KC1  +  MnCl2  +  5C1. 
Molecular:  2KMnO4  +  16HC1  ->  8H20  +  2KC1  +  2MnCl2  -f  5C12. 

The  combined  oxygen  of  the  permanganate  has  oxidized  the  hydro- 
gen chloride,  just  as  did  the  free  oxygen  in  Deacon's  process. 

Other  Means  of  Oxidizing  Hydrogen  Chloride.  —  Many 
other  compounds  of  oxygen  with  metals  interact  with  hydro- 
chloric acid  to  give  free  chlorine.  Lead  dioxide  Pb02,  potassium 
chlorate  KClOs,  potassium  dichromate  K2Cr20?,  and  manganese 
dioxide  MnO2,  are  of  this  nature.  The  last,  being  inexpensive,  is 
commonly  used  in  making  chlorine.  Being  an  insoluble  substance, 


158  COLLEGE    CHEMISTRY 

however,  the  manganese  dioxide  acts  much  more  slowly  than  does 
the  potassium  permanganate,  which  is  soluble.  A  large  amount  of 
the  materials,  and  the  aid  of  heat,  are  required  to  secure  a  rapid 
stream  of  chlorine. 

Manganese  Dioxide  and  Hydrogen  Chloride.  —  The  action 
of  manganese  dioxide  upon  hydrochloric  acid  is  an  instructive  one. 
It  is  a  general  rule,  of  which  we  shall  meet  many  applications,  that 
when  an  acid  interacts  with  an  oxide  of  a  metal,  there  are  two  con- 
stant features  in  the  result,  namely:  (1)  The  oxygen  of  the  oxide 
combines  with  the  hydrogen  of  the  acid  to  form  water,  and  (2)  the 
metal  of  the  oxide  combines  with  the  acid  radical  of  the  acid  accord- 
ing to  the  valences  of  each.  Here  the  skeleton  equation  should  be 
Mn02  +  HC1  ->  H2O  +  MnCLt.  With  O2,  to  form  water,  4HC1  is 
required,  and  the  product  is  2H2O.  Hence  the  equation  is 

Balanced:  Mn02  +  4HC1  ->  2H2O  +  MnCU. 

This  is  what  happens  in  the  first  place.  The  products  actually 
obtained,  however,  are  water,  manganous  chloride  MnCl2  and 
chlorine.  The  manganese  tetrachloride  can  be  preserved  by  cool- 
ing the  mixture.  It  is  decomposed  by  the  heating,  the  chlorine 
escapes,  and  the  other  two  products  remain  in  the  vessel. 

Mn02  -I-  4HC1  ->  2H20  +  MnCl2  +  C12.  (1) 

We  owe  the  chlorine  to  the  fact  that  the  tetrachloride  is  unstable. 
If  we  had  used  manganous  oxide  MnO,  we  should  have  had  a 
double  decomposition: 

MnO  +  2HC1  -*  H2O  +  MnCl2,  (2) 

but  we  should  have  got  no  chlorine.  Perhaps  the  simplest  way  to 
describe  the  difference  between  these  two  actions  is  in  terms  of  the 
valence  of  the  manganese.  In  MnIYO2n  the  element  is  quadriva- 
lent. This  means  that  its  atomic  weight  professes  to  be  able  to 
hold  four  atomic  weights  of  a  univalent  element.  The  four  valences 
of  oxygen  (20n)  can  do  the  same  thing.  In  equation  (1)  the 
oxygen  fulfils  this  promise  by  taking  4H1.  But  the  MnIV  can  hold 
only  2C11,  permanently,  and  lets  the  other  2C11  go  free.  In  other 
words,  the  valence  of  the  atomic  weight  of  manganese  changes  in  the 
course  of  the  action.  In  equation  (2),  on  the  other  hand,  the 
manganese  is  bivalent  to  start  with  (MnnOn),  and  is  able  to  retain 


CHLORINE  159 


the  amount  of  chlorine  (2C11)  equivalent  to  On.  Actions  like  that 
of  manganese  dioxide  in  (1)  are  classed  as  oxidations.  The  hydro- 
gen chloride,  or  rather  half  of  it,  is  oxidized.  A  graphic  mode  of 
writing  may  make  this  remark  clearer: 

,0  +  2HC1  -»  H2O  +  MnnCl2 


The  upper  half  is  a  double  decomposition,  the  lower  an  oxidation 
by  half  the  combined  oxygen  of  the  dioxide.  The  same  explana- 
tion applies  to  the  interaction  of  lead  dioxide  with  hydrochloric 
acid. 

Physical  Properties.  —  Chlorine  differs  from  the  gases  we 
have  encountered  so  far  in  having  a  strong  greenish-yellow  tint, 
a  fact  which  gave  rise  to  its  name  (Gk.,  pale  green),  and  having  a 
powerful,  irritating  effect  upon  the  membranes  of  the  nose  and 
throat. 

Density  (H  =  1),  35.79.  Boiling-point  (liq.),-33.6°. 

Weight  of  22.4  1.,  72.13  g.  Melting-point  (solid),  -102°. 

Sol'ty  in  Aq  (20°),  215  vols.  in  100.  Vap.  tension  (liq.)  0°,  3.66  atmos. 

Grit,  temp.,  +146°.  Vap.  tension  (liq.)  20°,  6.62  atmos. 

Since  the  G.M.V.  of  air  weighs  28.95  g.,  chlorine  is  two  and  a  half 
times  heavier.  In  solubility  it  stands  between  slightly  soluble 
gases,  like  oxygen  and  hydrogen,  and  those  which  are  extremely 
soluble.  It  can  be  collected  over  hot  water  or  a  strong  solution  of 
salt. 

Chlorine  was  first  liquefied  by  Northmore  (1806).  It  forms  a 
yellow  liquid  which,  contained  in  steel  cylinders  lined  with  lead, 
is  now  an  article  of  commerce.  On  being  cooled  below  —  102°,  it 
gives  a  pale-yellow  solid. 

In  recalling  the  physical  properties  of  a  gas,  remember  that  six 
(p.  31)  are  required:  color,  taste,  odor,  density,  solubility,  lique- 
fiability. 

Chemical  Properties.  —  Chlorine  is  at  least  as  active  a  sub- 
stance as  is  oxygen.  It  presents  a  more  varied  array  of  chemical 
properties  than  does  that  element.  The  binary  compounds  are 
called  chlorides. 


160  COLLEGE    CHEMISTRY 

Combines  with  Metals.  —  Powdered  antimony  (cold),  when 
thrown  into  chlorine,  unites  with  it  to  form  the  chloride  SbClg, 
which  appears  partly  as  vapor  and  partly  as  glowing  particles. 
Balanced:  Sb  +  3C1  -»  SbCl3. 

Molecular:  2Sb  +  3C12  ->  2SbCl3. 

Copper,  in  the  condition  of  thin  leaf  commonly  used  for  gilding 
(Dutch-metal),  catches  fire  when  thrust  into  the  gas,  giving  a  fog 
of  solid  cupric  chloride  CuCb  Sodium  burns  brilliantly,  giving 
a  cloud  of  sodium  chloride.  The  union  of  a  metal  like  sodium 
and  a  colored,  irritating  gas  to  give  a  mild  household  article,  like 
common  salt,  illustrates  the  extraordinary  nature  of  chemical 
change.  All  the  familiar  metals,  with  the  exceptions  of  gold  and 
platinum,  combine  readily  with  chlorine. 

When  metals  (like  copper  and  iron)  and  chlorine  are  first  thor- 
oughly freed  from  moisture,  combination  no  longer  occurs.  A 
trace  of  water  is  required  in  these,  as  it  is  in  many  other  chemical 
actions,  as  a  contact  agent.  Hence,  the  chlorine,  before  being 
compressed  into  steel  cylinders,  must  be  freed  entirely  from  water 
vapor  (see  Detinning). 

Combines  with  Hydrogen.  —  A  jet  of  hydrogen  burns 
vigorously  in  chlorine,  producing  hydrogen  chloride  HC1.  The 
union  of  the  gases,  when  a  mixture  of  them  is  kept  cold  and  in  the 
dark,  is  too  slow  to  be  perceived.  On  exposure  to  diffused  light, 
however,  they  unite  slowly,  while  a  sudden  flash  of  sunlight  or  the 
burning  of  a  magnesium  ribbon  causes  instant  explosion.  The 
effect  of  the  light  is  catalytic. 

Interacts    with     Compounds     Containing    Hydrogen. — 

When  a  lighted  taper  is  plunged  into  chlorine  it  continues  to  burn, 
but  a  dense  cloud  of  soot  (free  carbon)  rises  from  the  flame.  Blow- 
ing the  breath  into  the  jar  then  gives  the  fog  which  shows  the 
presence  of  hydrogen  chloride.  Thus  the  presence  of  hydrogen 
and  carbon  in  the  wax  is  proved.  We  learn,  also,  that  chlorine 
has  little  tendency  to  combine  with  carbon,  for  this  element  goes 
free.  A  few  drops  of  warm  turpentine,  poured  upon  a  strip  of 
paper,  when  placed  in  chlorine  give  a  violent  reaction  and  a  cloud 
of  finely  divided  carbon  bursts  forth. 

CioH16  +  8C12  ->  16HC1  +  IOC. 


CHLORINE  161 

Elements  Displaced  by  Chlorine.  —  The  action  on  turpen- 
tine is  a  displacement  of  the  carbon  by  the  chlorine.  Of  the  same 
nature  is  the  action  of  chlorine  upon  potassium  iodide  KI,  dry  or 
in  solution. 

C12-»2KC1  +  I2. 


The  iodine,  when  moist,  is  deep  brown  in  color.  A  mere  trace  of 
chlorine,  liberating  a  trace  of  iodine,  gives  no  visible  effect.  But 
if  some  starch  is  present,  even  a  trace  of  free  iodine  yields  a  deep 
blue  color.  This  reaction  is  used  as  a  test  for  chlorine,  for  free 
iodine  from  any  source,  and  for  starch  (p.  3).  To  test  for  chlorine, 
strips  of  filter  paper,  dipped  in  starch  emulsion  (starch  boiled  with 
much  water  and  cooled)  to  which  a  few  drops  of  potassium  iodide 
have  been  added,  are  used.  Combined  iodine,  as  in  potassium 
iodide,  has  no  effect  upon  starch.  Combined  chlorine,  as  in 
sodium  chloride,  has  no  action  upon  the  prepared  strips  of  paper  — 
free  chlorine  is  required. 

Action  Upon  Water.  —  We  have  seen  that  chlorine  seizes  the 
hydrogen  in  turpentine.  We  have  also  learned  that  it  combines 
with  the  hydrogen  in  steam,  reversing  Deacon's  process  to  the 
extent  of  20  per  cent.  It  also  acts  upon  cold  water,  when  dissolved 
in  the  latter,  although  in  a  similarly  incomplete  way.  The  sub- 
stances formed  are  hydrochloric  acid  and  hypochlorous  acid  HC10  : 

H2O  +  C12  ^  HC1  +  HC1O. 

With  half  -saturated  chlorine-water  at  10°  --that  is,  water  con- 
taining an  equal  volume  of  chlorine  gas  —  33  per  cent  of  the 
chlorine  is  changed  into  the  acids.  Thus,  chlorine-water  (the 
solution)  is  a  mixture  containing  dissolved  chlorine  and  two  acids. 
Hypochlorous  acid  (q.v.)  is  of  especial  interest  because  it  is  a  very 
active  substance,  with  powerful  oxidizing  qualities,  and  bleaches 
dyes  by  decomposing  them. 

The  action  comes  to  a  standstill  when  one-third  completed, 
because  the  two  acids  interact  to  reproduce  chlorine  and  water 
(read  the  equation  backwards).  The  action  is  reversible.  When 
the  solution  is  exposed  to  sunlight,  the  hypochlorous  acid  decom- 
poses and  oxygen  gas  is  liberated  and  escapes: 

2HC10  -*  2HC1  +  02  T  • 


162  COLLEGE    CHEMISTRY 

Since  this  removes  the  hypochlorous  acid,  on  whose  interaction 
with  the  hydrogen  chloride  the  reverse  action  depends,  the  for- 
ward action  proceeds  under  continuous  illumination  gradually  to 
completion.  Hence  the  aqueous  solution  of  chlorine  must  be  kept 
in  the  dark,  since  otherwise,  after  a  time,  a  dilute  solution  of 
hydrogen  chloride  alone  remains. 

The  reader  should  note  here  the  displacement  of  the  equilibrium, 
a  chemical  one  in  this  case,  in  consequence  of  the  annulment  of  one 
of  the  opposing  tendencies  (p.  90).  Through  the  destruction  of 
the  hypochlorous  acid,  one  of  the  tendencies,  namely  that  repre- 
sented in  the  backward  action,  becomes  inoperative.  The  for- 
ward action  is  not  itself  assisted,  but  it  is  no  longer  impeded,  and 
so  proceeds  to  completion. 

Action  by  Substitution.  —  When  actions  like  that  on  tur- 
pentine —  that  is  on  compounds  containing  carbon  and  hydrogen 
—  are  moderated  by  altering  the  conditions,  the  decomposition 
is  not  so  complete.  Using  a  lower  temperature  is  effective.  Thus, 
if  methane  CH*  (marsh-gas),  the  chief  component  of  natural  gas, 
is  mixed  with  chlorine  and  exposed  to  sunlight,  a  slower  action 
occurs,  of  which  the  first  stage  consists  in  the  removal  of  one  unit 
weight  of  hydrogen  and  the  substitution  of  chlorine  for  it  according 
to  the  following  equation: 

CH4  +  C12  -»  CH3C1  +  HC1. 

The  process  may  continue  further  by  the  substitution  *  of  chlorine 
for  the  units  of  hydrogen  one  by  one  until  carbon  tetrachloride 
CCU  is  finally  formed. 

The  action  on  water  is  a  substitution. 

Combines  with  Non-metals.  —  Phosphorus  burns  in  chlorine 
with  a  rather  feeble  light,  producing  primarily  phosphorus  tri- 

*  Substitution  resembles  displacement  (p.  55)  in  that  an  element  and  a 
compound  interact,  and  the  element  takes  the  place  of  one  unit  in  the  com- 
position of  the  latter.  In  the  above  action,  one  unit  of  chlorine  takes  the 
place  of  one  unit  of  hydrogen.  But  the  latter  is  not  liberated;  it  combines 
with  another  unit  of  chlorine.  The  action  resembles  double  decomposition, 
excepting  that  one  of  the  substances  is  not  a  compound,  but  a  diatomic  ele- 
ment. The  name  used  is  intended  to  fix  the  attention  on  the  compound  and 
on  the  fact  that  one  unit  has  been  substituted  for  another  in  it.  This  concep- 
tion is  a  favorite  one  in  the  chemistry  of  compounds  of  carbon. 


CHLORINE  163 

chloride  PC13,  a  liquid  (b.-p.  74°).  If  excess  of  chlorine  is  present, 
then,  as  the  trichloride  cools,  it  combines  to  form  the  solid  penta- 
chloride  PCls.  Sulphur,  when  heated,  unites  more  slowly,  giving 
sulphur  monochloride  S2C12,  a  liquid  used  in  vulcanizing  rubber. 
Chlorine  does  not  combine  directly  with  carbon,  nitrogen,  or  oxy- 
gen, although  compounds  with  those  elements  can  be  made  in- 
directly. With  the  helium  group  of  elements  (q.v.),  it  forms  no 
compounds. 

Combines  with  Compounds.  —  Chlorine  unites  with  many 
compounds.  Thus,  one  of  the  oxides  of  carbon,  carbon  monoxide 
CO,  when  mixed  with  chlorine  and  exposed  to  sunlight  gives  drops 
of  a  volatile  liquid  (b.-p.  8.2°)  known  as  phosgene  COC12. 

When  chlorine-water  is  cooled  with  ice,  a  compound,  chlorine 
hydrate  C12,8H2O  crystallizes  out.  Faraday  (1823)  placed  this 
substance  in  the  closed  limb  of  a  A-tube, 
sealed  the  open  end,  and  placed  the  empty 
limb  in  cold  water  (Fig.  68).  When  the 
hydrate  was  gently  warmed,  chlorine  gas 
was  given  off  and  was  liquefied  by  its  own 
pressure  in  the  cold  part  of  the  tube. 

FIG.  68. 

Chemical  Relations  of  the  Element.*  —  In  the  chlorides, 
an  atomic  weight  of  chlorine  is  equivalent  to  one  atomic  weight 
of  hydrogen  or  of  sodium.  The  element  is,  therefore,  univalent 
(p.  62).  It  never  shows  any  higher  valence  than  this,  save  in 
its  oxygen  compounds  (see  Chap.  XXIII).  The  oxides  of  chlo- 
rine interact  with  water  to  give  acids,  and  the  element  is,  there- 
fore, to  be  classed  as  a  non-metal  (p.  94).  It  belongs  to  that 
group  of  the  non-metals  called  the  halogens,  as  a  consideration  of 
some  others  of  its  relations  will  show  (see  Chap.  XV). 

*  In  accordance  with  the  distinction  that  must  be  drawn  (p.  16)  between 
the  element  as  a  variety  of  matter  in  combination,  and  the  elementary  sub- 
stance or  free  form  of  the  element,  and  to  avoid  a  common  source  of  con- 
fusion, we  shall  always  give  only  the  behavior  of  the  elementary  substance 
under  the  title  chemical  properties.  The  characteristics  which  distinguish 
the  compounds  of  the  element,  as  a  class,  from,  or  relate  them  as  a  class  to  the 
compounds  of  other  elements  will  then  appear  in  a  separate  section  u 
the  title  "  Chemical  relations"  (see  pp.  192,  208). 


164 


COLLEGE    CHEMISTRY 


Uses  of  Chlorine.  —  Large  quantities  of  chlorine  are  manu- 
factured for  the  preparation  of  bleaching  materials  and  disinfect- 
ing agents.  In  disinfection,  the  minute  germs  of  disease  and 
putrefaction  are  acted  upon  either  by  the  chlorine  or  by  the  hypo- 
chlorous  acid  formed  by  its  interaction  with  water,  and  instantly 
their  life  is  destroyed. 

Chlorides.  —  The  chlorides  are  described  individually  under 
the  other  element  which  each  contains.  The  majority  of  the 
chlorides  of  the  metals  are  easily  soluble  in  water.  The  chief 
exceptions  are  silver  chloride  AgCl,  mercurous  chloride  (calomel) 

HgCl,  cuprous  chloride  CuCl,  and  lead 
chloride  PbCl2.  The  last  of  these  is  on 
the  border  line  as  regards  solubility.  An 
appreciable  amount  dissolves  in  cold 
water,  and  a  considerable  amount  in 
boiling  water  (see  Table  of  Solubilities, 
inside  the  cover  at  the  front  of  this 
book).  For  the  various  modes  of  pre- 
paring chlorides  see  p.  146. 


Composition  of  Hydrogen  Chlo- 
ride.—  Being  now  familiar  with  both 
hydrogen  and  chlorine,  we  may  take  up 
the  question  of  the  proportion  by  vol- 
ume in  which  the  constituents  unite, 
and  the  relation  of  this  to  the  volume 
of  the  resulting  hydrogen  chloride. 

The  decomposition  of  the  solution  of 
hydrogen  chloride  in  water  by  means  of 
the  electric  current  proves  that  the  gases 
are  liberated  in  equal  volumes.  Brown- 
lee's  apparatus  for  demonstrating  this  is  shown  in  Fig.  69 .  The  cen- 
tral part  is  the  same  as  in  Fig.  27,  but,  when  the  three-way  stop- 
cock is  closed,  the  gases  go  to  right  and  left,  and  displace  the 
liquid  in  two  outside  tubes.  The  equal  rate  at  which  this  takes  place 
on  both  sides  proves  that  the  gases  are  generated  in  equal  volumes. 
In  order  to  ascertain  the  relation  between  the  volumes  of  the 
constituents  and  that  of  the  product,  we  may  unite  the  gases  and  find 


Fia.  69. 


CHLORINE 


165 


out  whether  any  change  in  volume  occurs.  A  tube  with  thick  walls 
(Fig.  70)  is  filled  with  the  mixed  gases  obtained  by  electrolysis. 
By  dipping  one  end  of  the  tube  under  mercury  and  . 

opening  the  lower  stopcock,  it  is  seen  that  no  gas  leaves        nT 
and  no  mercury  enters.    After  the  mixture  has  been        r  \ 
exploded,  by  the  light  from  burning  magnesium,  the 
same  test  is  repeated  with  the  same  result.     The  pres- 
sure has  therefore  remained  equal  to  that  of -the  at- 
mosphere.   Hence  there  has  been  no  change  in  volume 
as  the  result  of  the  union.     It  appears,  therefore,  that 

1  vol.  hydrogen  +  1  vol.  chlorine  — * 
2  vols.  hydrogen  chloride, 


a  result  in  harmony  with  Gay-Lussac's  law  (p.  98). 


FIG.  70. 


Confirmation  of  the  Formulse  C12  and  H2.  —  According  to 
Avogadro's  law,  there  are  equal  numbers  of  molecules  in  equal 
volumes  of  these  gases.  When  hydrogen  and  chlorine  combine, 
one  volume  of  each  of  these  gases  gives  two  volumes  of  hydrogen 
chloride.  Let  us  imagine  the  experiment  to  be  made  with  minute 
volumes  holding  one  hundred  molecules  each: 

HYDROGEN  CHLORIDE  HYDROGEN          CHLORINE 

came  from 


100 

100 

The  200  molecules  of  hydrogen  chloride  must  contain  at  least  200 
fragments  of  chlorine,  since  there  is  a  sample  in  each  molecule. 
Now  the  200  fragments  of  chlorine  came  from  a  volume  contain- 
ing only  100  molecules  of  chlorine.  Each  of  these  must  therefore 
have  been  split  in  the  chemical  action.  The  same  is  true  of  each 
molecule  of  hydrogen.  Hence  the  molecules  of  free  hydrogen 
and  free  chlorine  contain  at  least  two  atoms.  If  we  consider  the 
molecular  formula  of  a  substance  as  representing  one  molecule, 
the  equation  for  this  action  is: 

H2  +  Cla  ->  2HC1. 

There  are  two  molecules  on  each  side  of  the  equation,  and  this 
corresponds  with  the  fact  that  there  is  no  change  in  the  total 
volume. 


166  COLLEGE    CHEMISTRY 

Classification  of  Chemical  Interactions  and  Exercises 
Thereon.  —  So  far  we  have  defined  ten  more  or  less  distinct  kinds 
of  chemical  change,  seven  differing  in  mechanism:  Combination 
(p.  7),  decomposition  (p.  14),  dissociation  (p.  93),  displacement  (p. 
55),  substitution  (p.  162),  double  decomposition  (pp.  142,  147),  and 
internal  rearrangement  (p.  148) ;  and  three  others :  oxidation  (pp.  36, 
156,  158),  reduction  (pp.  37,  59),  and  electrolysis  (pp.  55,  155). 
Illustrations  of  all  but  one  of  these  will  be  found  in  the  present 
chapter.  Some  actions  belong  to  one  of  the  first  seven,  and  also 
to  one  of  the  three  other  classes.  The  ability  readily  to  classify 
each  phenomenon,  as  it  comes  up,  requires  precisely  that  grasp 
of  the  framework  of  the  science  which  the  reader  must  seek  speedily 
to  attain.  ^ Fpr^examplc,  let  him  classjfy-  the  folloAving  actions:  1. 

>f  nea 


action  of  potassium  on  water;  2.  of  neat,  on  i>ptassiumJehlorate; 
3.  pf  chlorine  on  metals;  4.  of  chlorine  on  turpentine;  5.  of 
chlorine  on  potassium  iodide;  6.  of  chlorine  on  methane;  7.  of 
carbon  monpxidi^^c^/c^orin^j^  8.  of  sunlight  on  hypochlorous 
acid;  9.  or  sulphuric  acia  on  sll£;  10.  of  zinc  oxide  and  hydro- 
chloric acid;  11.  of  zinc  on  hydrochloric  acid. 

12.  In  the  interactions  of  potassium  permanganate  and  of  man- 
ganese dioxide,  respectively,  with  hydrochloric  acid,  what  fractions 
of  the  whole  chlorine  are  liberated?    What  are  the  commercial 
advantages  of  the  use  of  salt  and  sulphuric  acid  with  the  manganese 
dioxide? 

13.  In  view  of  the  explanations  given,  define  the  general  nature 
of  the  substances  (p.  157)  which  may  be  used  to  oxidize  hydro- 
chloric acid. 

14.  What  are  the  relative  volumes  of  the  gaseous  interacting 
substances  and  products  in  the  following  reactions :   (a)  turpentine 
vapor  and  chlorine;    (b)  methane  and  chlorine;    (c)  phosphorus 
vapor  and  chlorine;   (d)  carbon  monoxide  and  chlorine. 


CHAPTER  XIII 
ENERGY  AND   CHEMICAL   CHANGE 

IN  describing  chemical  changes,  the  fact  that  heat  was  evolved 
has  frequently  been  mentioned.  In  several  instances'a  current  of 
electricity  has  been  used  to  produce  chemical  change.  It  is  now 
necessary  to  collect  these  scattered  facts  and  classify  them  for 
future  use. 

Physical  Accompaniments  of  Chemical  Change. — When 
iron  and  sulphur  combined  (p.  13),  and  when  iron  burned  in  oxy- 
gen or  copper  in  chlorine,  much  heat  was  developed.  On  the 
other  hand,  the  decomposition  of  mercuric  oxide,  as  was  pointed 
out  (p.  14),  owed  its  continuance  to  the  persistent  application  of 
heat  and  ceased  as  soon  as  the  source  of  heat  was  withdrawn. 
Here,  apparently,  heat  was  consumed  during  the  progress  of  the 
change,  and  the  chemical  action  was  limited  by  the  amount  of 
heat  supplied.  The  production  or  consumption  of  heat  may,  there- 
fore, be  a  feature  of  chemical  change. 

In  the  burning  of  iron  or  magnesium  in  oxygen,  and  in  the 
actions  of  chlorine  on  copper  and  turpentine,  light  was  also  pro- 
duced. Conversely,  silver  chloride  (p.  147)  can  be  kept  any  length 
of  time  in  the  dark,  but  in  sunlight  it  becomes  first  bluish  and 
then  brown,  simultaneously  giving  off  chlorine  gas  and  finally 
leaving  only  silver  as  a  fine  powder.  Silver  bromide  or  iodide,  in 
photographic  plates,  films,  and  paper,  is  changed  by  light  in  a 
similar  way,  liberating  the  bromine  or  iodine.  It  would  appear, 
therefore,  that  light  may  be  given  out  or  consumed  in  connection 
with  chemical  change. 

We  have  seen  (p.  155)  that  a  current  of  electricity  may  be  em- 
ployed to  decompose  hydrochloric  acid  and  other  chlorides,  and 
the  battery,  or  other  source  of  the  current  must  be  kept  going  or 
the  chemical  change  stops.  The  inverse  of  this  is  likewise  familiar. 
If  we  place  in  dilute  sulphuric  acid  a  stick  of  the  metal  zinc,  we 

167 


168 


COLLEGE    CHEMISTRY 


\  v^7.:,>'> 


find  that  hydrogen  is  given  off  (Fig.  71),  that  the  zinc  goes  into 
solution  as  zinc  sulphate  (p.  53),  and  that  a  large  amount  of  heat 
is  developed.  If  zinc  in  fine  particles,  with 
much  surface,  is  used,  the  liquid  may  even 
rise  spontaneously  to  the  boiling-point. 
This  form  of  the  action  produces  heat.  If, 
however,  we  attach  the  same  stick  of  zinc 
to  a  copper  wire  and,  having  provided  a 
plate  of  platinum  also  connected  with  a  wire, 
immerse  the  two  simultaneously  in  the  acid 
(Fig.  72),  then  a  galvanometer,  with  which 
the  wires  are  connected,  shows  at  once  the 
passage  of  a  current  of  electricity  round 
the  circuit.  Exactly  the  same  chemical 
change  goes  on  as  before.  The  sole  differ- 
ence is  that  the  gas  appears  to  arise  from 
the  surface  of  the  platinum.  It  is  easy  to 
show,  however,  that  the  platinum  by  itself 
FlG-  7L  is  not  acted  upon  by  dilute  acids  and,  in 

this  case,  undergoes  no  change  whatever;  it  serves  simply  as  a 
suitable  conductor  for  the  electricity.     Here,  then,  in  place  of  the 


FIG.  72. 


heat  which  the  first  plan  produced,  we  get  an  electric  current. 
The  arrangement  is,  in  fact,  a  battery-cell,  for  a  battery  is  a 
system  in  which  a  chemical  action  which  would  otherwise  give 


ENERGY  AND    CHEMICAL   CHANGE  169 

heat  furnishes  electricity  instead.  Thus,  electrical  energy  may  be 
consumed  or  produced  in  connection  with  a  change  in  composition. 

Even  violent  rubbing  in  a  mortar,  in  the  case  of  some  substances, 
can  effect  an  appreciable  amount  of  decomposition  in  a  few  min- 
utes. In  this  way  silver  chloride  can  be  separated  into  silver  and 
chlorine,  just  as  by  light.  It  is  the  mechanical  energy  which  is  the 
agent,  and  part  of  it  is  consumed  in  producing  the  change,  and 
only  the  balance  appears  as  heat.  Conversely,  the  production  of 
mechanical  energy,  as  the  result  of  chemical  change,  is  seen  in 
the  behavior  of  explosives  and  in  the  working  of  our  muscles. 
Thus,  mechanical  energy  may  be  used  up  or  produced  in  chemical 
changes. 

Summing  up  our  experience,  we  may  state  that  no  change  in 
composition  occurs  without  some  accompaniments,  such  as  the 
production  or  consumption  of  heat,  light,  electrical  energy,  or,  in 
some  cases,  mechanical  energy. 

Classification  of  the  Accompaniments  of  Change  in  Com- 
position: Energy.  —  The  problem  of  classifying  (i.e.,  placing  in  a 
suitable  category)  things  like  heat,  light,  and  electricity  has  occu- 
pied much  attention.  In  all  changes  in  composition,  one  of  these 
natural  accompaniments  is  given  out  or  absorbed,  sometimes  in 
great  amount,  yet  in  none  is  any  alteration  in  weight  observed.* 
There  are  many  things  which  are  real,  however,  even  if  they  are 
not  affected  by  gravitation.  In  the  present  instance  we  reason 
as  follows: 

A  brick  in  motion  is  different  from  a  brick  at  rest.  The  former 
can  do  some  things  that  the  latter  cannot.  Furthermore,  we  can 
easily  make  a  distinction  in  our  minds.  The  brick  can  be  deprived 
of  the  motion  and  be  endowed  with  it  again.  Thus,  we  can  get 
the  idea  of  motion  as  a  separate  conception.  Similarly,  we  observe 
that  a  piece  of  iron  behaves  differently  when  hot,  and  when  cold, 
when  bearing  a  current  of  electricity,  and  when  bearing  none. 
We  conceive  then  of  the  brick  or  the  iron  as  having  a  certain 
amount  and  kind  of  matter  which  is  unalterable,  and  as  having 
motion,  heat,  or  electricity  added  to  this  or  removed.  Thus,  we 
describe  our  observations  by  using  two  categories,  one  of  which 

*  Electrons  (q.v.}  do  possess  mass,  but  it  is  very  small  compared  with  that 
of  the  materials  concerned. 


170 


COLLEGE    CHEMISTRY 


includes  the  various  kinds  of  matter,  and  the  other,  various  things 
whose  association  with  matter  seems  to  be  invariable  and  is  often 
so  conspicuous.  The  latter  we  call  the  forms  of  energy. 

The  Practical  Importance  of  Energy  in   Chemistry.  — 

The  absorption  or  liberation  of  energy  accompanying  a  chemical 
transformation  of  matter  is  often,  of  the  two,  the  more  important 
feature.  We  do  not  burn  coal  in  order  to  manufacture  carbon 
dioxide  gas.  We  are  glad  to  get  rid  of  the  material  product  through 
,the  chimney.  It  is  the  heat  we  want.  We  do  not  buy  gasoline 
(petrol)  for  an  automobile  in  order  to  obtain  various  gases  to  ex- 
pel through  the  muffler.  We  really  pay  for  the  mechanical  energy. 
It  is  the  same  with  burning  illuminating-gas  or  magnesium  powder 
when  we  want  light,  and  with  eating  food,  which  we  do,  chiefly,  to 
get  energy  to  sustain  our  activity.  We  do  not  run  electricity  for 
hours  into  a  storage  battery  in  order  to  make  a  particular  compound 

(lead  dioxide,  for  example),  but  in  order 
to  save  and  store  the  energy  for  future 
use.  In  industry  and  life  fully  half  the 
total  amount  of  chemical  change  in- 
volved is  set  in  motion  by  us,  solely  on 
account  of  the  energy  changes  it  in- 
volves. But  the  production  of  energy 
in  chemical  change  is  not  only  thus  of 
practical  importance;  it  is  also  of  scien- 
tific interest,  as  will  be  seen  in  the  sec- 
tion on  energy  and  chemical  activity 
(see  below). 


FIG.  73. 


Interconvertibility  of  Forms  of 
Energy:  Conservation.  —  At  first 
sight,  these  accompaniments  of  matter 
seem  to  be  quite  unrelated.  But  a  relation  between  them  can  be 
found.  If  the  heat  of  a  Bunsen  flame  or  of  the  sun  is  brought 
under  a  hot-air  motor  (Fig.  73)  violent  motion  results.  Again,  if 
the  motor  is  connected  with  a  dynamo,  electricity  may  be  generated. 
Still  again,  if  the  current  from  the  dynamo  flows  through  an  incan- 
descent lamp,  heat  and  light  are  evolved.  Conversely,  when  motion 
of  the  hot-air  motor  is  impeded  by  a  brake,  heat  appears.  When 


ENERGY  AND  CHEMICAL  CHANGE          171 

a  current  of  electricity  is  run  through  the  dynamo,  the  armature 
of  the  latter  turns  and  motion  results.  But  the  most  significant 
facts  are  still  to  be  mentioned.  The  heat  absorbed  by  the  motor 
is  found  to  be  greater  when  the  machine  is  permitted  to  move  and 
do  work,  than  when  it  is  not.  Thus,  it  is  found  that  when  work  is 
done  some  heat  disappears,  and  this  heat  is,  in  fact,  transformed 
into  work.  Similarly,  when  the  poles  of  the  dynamo  are  properly 
connected  and  electricity  is  being  produced,  and  only  then,  motion 
is  used  up.  This  is  shown  by  the  effort  required  to  turn  the  arma- 
ture under  these  circumstances,  and  the  ease  with  which  it  is 
turned  when  the  circuit  is  open.  So,  with  a  conductor  like  the 
filament  in  the  lamp,  unless  it  offers  resistance  to  the  current  and 
destroys  a  sufficient  amount  of  electrical  energy,  it  gives  out  neither 
light  nor  heat.  Finally,  motion  gives  no  heat  unless  the  brake  is 
set,  and  effort  is  then  demanded  to  maintain  the  motion.  These 
experiences  lead  us  to  believe  that  we  have  here  a  set  of  things 
which  are  fundamentally  of  the  same  kind,  for  each  form  can  be 
made  from  any  of  the  others.  We  have,  therefore,  invented  the 
conception  of  a  single  thing,  of  which  heat,  light,  electricity,  and 
motion  are  forms,  and  to  it  we  give  the  name  energy:  energy  is 
work  and  every  other  thing  which  can  arise  from  work  and  be  con- 
verted into  work  (Ostwald). 

Closer  study  shows'  that  equal  amounts  of  electrical  or  mechani- 
cal energy  always  produce  equal  amounts  of  heat.  No  loss  is 
ever  observed  in  the  transformations  of  energy,  any  more  than 
in  the  transformations  of  matter.  Hence,  J.  R.  Mayer  (1842), 
Colding  (1843),  and  Helmholtz  (1847)  were  led  independently  to 
the  conclusion  that  in  a  limited  system  no  gain  or  loss  of  energy 
is  ever  observed.  This  brief  statement  of  the  results  of  many  ex- 
periments is  called  the  law  of  the  conservation  of  energy. 

Application  of  the  Conception  of  Energy  in  Chemistry.  - 

At  first  sight  it  looks  as  if  the  statement  that  energy  is  conserved 
is  not  applicable  in  chemistry.  Heat  and  electricity,  for  example, 
seem  to  be  produced  and  consumed,  in  connection  with  changes  in 
composition,  in  a  mysterious  manner.  We  trace  light  in  an  in- 
candescent lamp  back  to  the  electricity,  and  this  in  turn  to  the 
mechanical  energy,  and  this  again  to  the  heat  in  the  engine.  But 
what  form  of  energy  gave  the  heat  developed  by  the  combustion 


172  COLLEGE    CHEMISTRY 

of  the  coal  under  the  boiler,  or  by  the  union  of  iron  and  sulphur 
in  our  first  experiment?  Since  we  do  not  perceive  any  electricity, 
light,  heat,  or  motion,  in  the  original  materials,  and  yet  wish  to  create 
an  harmonious  system,  we  are  bound  to  conceive  of  the  iron  and  the 
sulphur,  and  the  coal  and  the  air,  as  containing  another  form  of 
energy,  which  we  call  internal  energy.  Similarly,  when  heat  is  used 
up  in  decomposing  mercuric  oxide,  or  light  in  decomposing  silver 
chloride,  we  regard  the  energy  as  passing  into,  and  being  stored 
in  the  products  of  decomposition  in  the  form  of  internal  energy. 

These  conclusions  compel  us,  for  the  sake  of  consistency,  to 
think  of  all  our  materials  as  repositories  of  energy  as  well  as  of 
matter,  each  of  these  two  constituents  being  equally  real  and 
equally  important.  A  piece  of  the  substance  known  as  "iron" 
must  thus  be  held  to  contain  so  much  iron  matter  and  so  much 
internal  energy.  So  ferrous  sulphide  contains  sulphur  matter, 
iron  matter,  and  internal  energy.  Thus,  by  a  substance  we  mean 
a  distinct  species  of  matter,  simple  or  compound,  with  its  appro- 
priate proportion  of  internal  energy. 

In  the  course  of  this  discussion  it  has  become  clear  that  it  is 
characteristic  of  chemical  phenomena  that,  besides  a  change  in  the 
nature  of  the  matter,  there  is  always  an  alteration  in  the  amount  of 
internal  energy  in  the  system.  This  alteration  involves  the  produc- 
tion of  internal  energy  from,  or  the  transformation  of  internal 
energy  into  some  other  form  of  energy. 

Energy  and  Chemical  Activity.  —  Other  things  being  equal, 
actions  in  which  there  is  a  relatively  large  loss  of  internal  energy 
proceed  rapidly;  that  is  to  say,  in  them  a  large  proportion  of  the 
material  is  changed  in  the  unit  of  time.  Those  in  which  less  en- 
ergy is  transformed  proceed  more  slowly.  The  speed  of  the  chemi- 
cal change,  and  the  quantity  of  energy  available  because  of  it,  are 
closely  related.  Now,  we  are  accustomed  to  speak  of  materials 
which,  like  iron  and  sulphur,  interact  rapidly  and  with  liberation 
of  much  energy  as  "chemically  active."  Thus,  relative  chemical 
activity  may  be  estimated,  (1)  by  observing  the  speed  of  a  change 
(see  below),  or,  in  many  cases  (2)  by  measuring  the  heat  developed 
(see  Thermo-chemistry,  below),  or  (3)  by  ascertaining  the  electro- 
motive force  of  the  current,  when  the  materials  are  arranged  in 
the  form  of  a  battery-cell  (see  Chap.  XXXIX). 


ENERGY  AND  CHEMICAL  CHANGE          173 

It  is  evident  that  the  chemical  activity  of  a  given  substance  will 
not  be  the  same  towards  all  others.  Thus,  iron  unites  much  more 
vigorously  with  chlorine  than  with  sulphur  and,  with  identical 
amounts  of  iron,  more  heat  is  liberated  in  the  former  case  than  in 
the  latter.  With  silver,  sodium,  and  many  other  substances,  iron 
does  not  unite  at  all.  One  of  the  tasks  of  the  chemist  is  to  make 
such  comparisons  as  this.  He  calls  the  results  the  specific  chemical 
properties  of  the  substances  in  question. 

The  i(  Cause  "   of  Chemical  Activity  or  Affinity.  —  The 

reader  will  undoubtedly  be  inclined  to  inquire  whether  we  can 
assign  any  cause  for  the  tendency  which  substances  have  to  undergo 
chemical  change.  Why  do  iron  and  sulphur  unite  to  form  ferrous 
sulphide,  while  other  pairs  of  elements  taken  at  random  will  fre- 
quently be  found  to  have  no  effect  upon  one  another  under  any 
circumstances?  The  answer  is  that  we  do  not  know.  Questions 
like  this  have  to  go  without  answer  in  all  sciences.  What  is  the 
cause  of  gravitation?  We  know  the  facts  which  are  associated 
with  the  word  —  the  fact  that  bodies  fall  towards  the  earth,  for 
example  —  but  why  they  fall  we  are  unable  to  say.  So,  with 
chemical  change,  we  can  state  all  the  facts  we  know  about  it,  but 
even  then  we  cannot  say  why  it  takes  place. 

The  word  cause  was  employed  in  the  heading  of  this  section,  and 
it  will  be  observed  that  no  cause  was  found.  This  is  the  invariable 
rule  in  physical  or  chemical  phenomena.  We  know  of  no  causes, 
in  the  sense  in  which  the  word  is  commonly  employed. 

The  word  cause  has  only  one  definite  use  in  science.  When  we 
find  that  thorough  incorporation  of  the  three  materials  is  needed  to 
secure  good  gunpowder,  we  say  that  the  intimate  mixing  is  a  cause 
of  its  being  highly  explosive.  By  this  we  simply  mean  that  intimate 
mixture  is  a  necessary  antecedent  of  the  result.  A  cause  is  a  condition 
or  occurrence  which  always  precedes  another  condition  or  occurrence. 

The  Speed  of  Chemical  Actions:  a  Means  of  Measuring 
Activity.  —  One  means  of  measuring  the  relative  chemical  activi- 
ties of  several  substances  is  to  observe  the  speed  with  which  they 
undergo  the  same  chemical  change.  Thus  we  may  compare  the 
activities  of  the  various  metals  by  allowing  them  separately  to 
interact  with  hydrochloric  acid  and  collecting  and  measuring  the 


174  COLLEGE    CHEMISTRY 

hydrogen  liberated  per  minute  by  each.  It  will  be  seen,  even  in  the 
roughest  experiment,  that  magnesium  is  thus  much  more  active 
than  zinc.  The  comparison  must  be  made  with  such  precautions, 
however,  as  will  make  it  certain  that  the  conditions  under  which 
the  several  metals  act  are  all  alike.  Thus,  in  spite  of  the  heat 
evolved  by  the  action,  means  must  be  used,  by  suitable  cooling,  to 
keep  the  temperature  at  some  fixed  point  during  the  experiment, 
for  all  actions  become  more  rapid  when  the  temperature  rises 
(p.  20).  Again,  the  pieces  of  the  various  metals  must  be  arranged 
so  that  equal  surfaces  are  exposed  to  the  acid  in  each  case.  It  is 
found  that  the  order  in  which  this  comparison  places  the  metals  is 
much  the  same  as  that  in  which  they  are  placed  by  a  study  of 
other  similar  actions.  A  single  table,  showing  the  order  of  activ- 
ity (p.  60),  suffices,  therefore,  for  all  purposes. 

Thermochemistry.  —  Chemical  changes  in  which  heat  is 
liberated  are  called  exothermal.  Those  in  which  heat  is  continu- 
ously absorbed  (pp.  14,  167,  113)  are  called  endothermal  changes. 
Since  the  activities,  or  affinities  of  two  substances  (say,  two  metals) 
may  often  be  measured  by  observing  the  amounts  of  heat  liberated 
when  each  combines  with  a  third  substance  (say,  oxygen),  it  will 
be  instructive  now  to  consider  some  of  the  elementary  facts  of 
thermochemistry. 

The  chemical  interactions  to  be  studied  thermally  are  arranged 
so  that  they  may  be  carried  out  in  a  small  vessel  which  can  be 
placed  inside  another  containing  water.  The  whole  apparatus  is 
called  a  calorimeter  (Gk.,  heat-measurer).  The  heat  developed 
raises  the  temperature  of  this  water.  Where  gases  like  oxygen 
are  concerned,  a  closed  bulb  of  platinum  forms  the  inner  vessel. 
The  quantity  of  heat  capable  of  raising  one  gram  of  water  one 
degree  in  temperature  at  15°  Centigrade  is  called  a  calorie.  So  that 
250  grams  of  water  raised  1°  would  represent  250  calories,  and  20 
grams  of  water  raised  5°  would  represent  100  calories. 

Thermochemical  Equations.  —  While  in  physics  the  unit 
of  quantity  is  the  gram,  in  chemistry  the  unit  which  we  select  is 
naturally  a  gram-atomic  weight  or  a  gram-molecular  weight  of  the 
substance.  Thus,  the  heat  of  combustion  of  carbon  means  the 
heat  produced  by  combining  twelve  grams  of  carbon  with  thirty- 


ENERGY  AND  CHEMICAL  CHANGE          175 

two  grams  of  oxygen,  and  is  sufficient  to  raise  nearly  100,000 
grams  of  water  one  degree.     This  is  expressed  as  follows: 
C  +  02  -*  C02  +  96,820  cal. 

In  other  words,  the  combustion  of  less  than  half  an  ounce  of  carbon 
will  raise  over  two  pounds  (one  kilogram)  of  water  from  0°  to  the 
boiling-point. 

When  the  action  is  one  which  absorbs  heat,  this  fact  is  indicated 
by  the  negative  sign  preceding  the  number  of  calories.     Thus,  the 
dissociation  of  36  g.  of  water  vapor  into  hydrogen  and  oxygen 
absorbs  28,800  cal.  per  gram  of  hydrogen  liberated: 
2H20  ->  2H2  +  02  -  115,200  cal. 

If  the  action  is  reversible,  as  this  one  is,  the  heat  absorbed  when 
it  proceeds  in  one  direction  is  equal  to  that  liberated  when  it  goes 
in  the  other  direction: 

2H2  +  02  ->  2H20  +  115,200  cal. 

An  action  which  absorbs  heat  can  take  place  only  when  heat  or 
some  other  form  of  energy  is  furnished.  Thus,  the  electrolysis  of 
aqueous  hydrochloric  acid  (p.  155)  consumes  electrical  energy, 
which  is  equivalent  in  amount  to  the  heat  given  out  when  hydrogen 
and  chlorine  unite  to  form  hydrogen  chloride,  plus  the  heat  liber- 
ated when  the  latter  dissolves: 

H2  +  C12  +  Aq  ->  2HC1,  Aq  +  78,800  cal. 

Answers  to  Possible  Questions.  —  It  is  always  found  that 
the  same  quantities  of  any  given  chemical  substances,  undergoing 
the  same  chemical  change  under  the  same  conditions,  produce 
or  absorb,  according  as  the  action  is  exothermal  or  endothermal, 
amounts  of  heat  which  are  equal. 

The  rate  at  which  a  given  chemical  action  is  allowed  to  take  place 
has  no  influence  on  the  total  amount  of  heat  consumed  or  produced. 
It  may  not  at  first  sight  appear  obvious  that  rusting  evolves  heat, 
but  a  delicate  thermometer  will  show  that  a  heap  of  rusting  nails  is 
somewhat  higher  in  temperature  than  surrounding  bodies.  Poor 
conductors,  like  oily  rags  and  ill-dried  hay,  show  a  tendency  to 
spontaneous  combustion  owing  to  accumulation  of  the  slowly 
developing  heat  of  oxidation  (p.  37).  The  warmth  of  our  own 
bodies  is  due  to  the  same  cause. 


176  COLLEGE    CHEMISTRY- 

It  should  be  noted  that  production  or^bsprption  of  heat  is  not, 
in  itself,  an  evidence  of  chemical  action.  Physical  changes  are  all 
likewise  accompanied  by  the  same  phenomena.  Thus,  the  evapo- 
ration of  water  absorbs  heat,  and  condensation  of  a  vapor  and  the 
crystallization  of  a  supercooled  liquid  liberate  heat. 

The  heat  of  solution  (cf.  pp.  125,  145)  is  the  heat  liberated  (or 
absorbed)  on  dissolving  one  mole  of  the  substance  in  a  large 
amount  of  water.  A  part  of  the  water  always  undergoes  chemical 
change  (p.  139) .  The  solute  also  frequently  combines  with  a  part 
of  the  water,  or  is  ionized  (q.v.\  and  the  change  in  volume  of  the 
mixture  (p.  138),  as  a  physical  phenomenon,  would  alone  entail  a 
heat-change.  Hence  this  heat  effect  is  partly  chemical  and  partly 
physical  in  origin. 

Exercises.  —  1.  Which  form  of  energy  is  delivered  as  such, 
and  paid  for  as  such,  in  most  cities? 

2.  How  many  calories  are  required  to  raise  500  g.  of  a  substance 
of  specific  heat  0.5  from  15°  to  37°  (p.  174)? 

3.  The  combustion  of  1  g.  of  sulphur  to  sulphur  dioxide  develops 
2220  calories.    What  is  the  heat  of  combustion  of  sulphur?    Write 
the  thermochemical  equation. 


CHAPTER  XIV 
CHEMICAL  EQUILIBRIUM 

IN  spite  of  its  formidable  title,  this  chapter  will  introduce  nothing 
novel.  Its  purpose  is  to  collect  together  and  organize  more 
definitely  a  number  of  scattered  facts  and  ideas  which  have  already 
come  up  in  various  connections.  On  this  account,  however,  it  will 
be  all  the  more  necessary  for  the  reader  to  refresh  his  remembrance 
of  these  facts  and  ideas  by  re-reading  all  pages  to  which  reference 
is  made. 

Reversible  Actions.  —  In  discussing  Deacon's  process  (p.  156), 
it  was  stated  that  the  action  comes  to  rest  although  a  large  amount 
of  both  of  the  interacting  substances  (20  per  cent  at  345°)  still 
remains  available:  (20  per  cent)  4HC1  +  02  <=*  2H20  +  2C12 
(80  per  cent).  Now  the  materials  thus  left  unused  are  presumably 
no  less  capable  of  interacting  than  were  the  parts  which  have 
already  reacted.  The  solution  of  this  mystery  lies  in  the  fact 
(p.  156)  that  the  products  themselves  interact  to  reproduce  the 
initial  substances  (read  the  equation  backwards).  Thus  two 
changes,  one  of  which  undoes  the  work  of  the  other,  are  going  on 
simultaneously.  In  consequence  of  this,  neither  action  can  reach 
completion.  As  we  should  expect,  experiment  shows  that  it 
makes  no  difference  whether  we  start  with  pure  chlorine  and  steam, 
or  with  hydrogen  chloride  and  oxygen;  the  proportions  of  the  four 
substances  found  in  the  tube,  after  it  has  been  kept  at  345°  for  a 
sufficient  time,  are  in  both  cases  the  same.  A  general  statement 
may  be  founded  on  facts  like  this,  to  the  effect  that  a  chemical 
action  must  remain  more  or  less  incomplete  when  the  reverse  action 
also  takes  place  under  the  same  conditions.  Two  arrows  pointing 
in  opposite  directions  are  used  in  equations  representing  reversible 
changes.* 

*  The  reader  must  avoid  the  idea  that  a  reversible  action  is  one  which  goes 
to  completion,  and  then  runs  back  to  a  certain  extent.  This  conception  would 
be  contrary  to  the  fact,  and  inexplicable  by  the  kinetic  method. 

177 


178  COLLEGE    CHEMISTRY 

The  foregoing  example  of  a  reversible  action,  and  the  following 
examples  which  very  closely  resemble  it,  should  now  be  looked  up 
and  studied  attentively.  The  discussion  in  this  and  the  following 
sections,  for  which  they  furnish  the  basis,  cannot  otherwise  be 
understood:  (1)  the  interaction  of  chlorine  and  water  (p.  161), 
which  was  fully  discussed  at  the  time;  (2)  the  behavior  of  phos- 
phorus pentachloride  vapor  (p.  117);  (3)  the  behavior  of  water 
vapor  (p.  93),  of  phosphorus  vapor  (p.  117),  of  sulphur  vapor 
(p.  117),  and  of  iodine  vapor  (p.  117). 

When  the  action  is  one  which  is  reversible,  but,  under  the  cir- 
cumstances being  discussed,  proceeds  farther  towards  completion 
in  one  direction  than  in  the  other,  the  arrow  will  be  modified  to 
indicate  this  fact: 

C12  +  H20 ±=?  HC1  +  HC10  (p.  161). 

When  this  relative  completeness  is  due  to  precipitation  or  vola- 
tilization, the  fact  may  be  indicated  by  vertical  arrows: 

NaCl  -f  H2S04i=>NaHS04  +  HC1T  (p.  141). 
NaClj  +  H2S04^NaHS04  +  HC1  (p.  143). 

Actions  Which  Proceed  to  Completion.  —  All  chemical 
actions  do  not  belong  to  the  reversible,  incomplete  class.  Many 
proceed  uninterruptedly  to  exhaustion  of  one,  or  all,  of  the  in- 
gredients. For  example,  equivalent  amounts  of  magnesium  and 
oxygen  combine  completely,  2Mg  -f-  02  — >  2MgO.  Here,  how- 
ever, the  product  is  not  decomposed  even  at  the  white  heat  pro- 
duced by  the  vigor  of  the  union.  Indeed,  magnesium  oxide  cannot 
be  decomposed,  and  the  action  reversed,  at  any  temperature  we 
can  command.  The  other  complete  actions,  like  the  decomposi- 
tion of  potassium  chlorate  (p.  27),  are  so  because  they  are  likewise 
irreversible. 

Explanation  in  Terms  of  Molecules.  —  Restating  these  facts 
in  terms  of  the  molecules  will  enable  us  to  reason  more  clearly 
about  this  variety  of  chemical  change.  Suppose  we  start  with  the 
materials  represented  on  one  side  only  of  such  an  equation,  say  the 
hydrogen  chloride  and  oxygen  in  that  on  p.  177.  The  molecules 
of  these  materials  will  encounter  one  another  frequently  in  the 
course  of  their  movements.  In  a  certain  proportion  of  these 


CHEMICAL   EQUILIBRIUM  179 

collisions  the  chemical  change  will  take  place.  In  the  earliest 
stages  there  will  be  few  of  the  new  kind  of  molecules  (say,  of 
chlorine  and  steam),  but,  as  the  action  goes  on,  these  will  increase 
in  number.  There  will  be  two  consequences  of  this.  In  the  first 
place  the  parent  materials  (in  this  case,  hydrogen  chloride  and 
oxygen)  will  diminish  in  amount,  the  collisions  between  their 
molecules  will  become  fewer,  and  the  speed  of  the  forward  action 
will  therefore  become  less  and  less.  In  the  second  place  the  in- 
crease in  the  number  of  molecules  of  the  products  will  result  hi 
more  frequent  collisions  between'them,  in  more  frequent  occurrence 
of  the  chemical  change  which  they  can  undergo,  and  thus  in  an 
increase  in  the  speed  of  the  reverse  action.  The  forward  action 
begins  at  its  maximum  and  decreases  in  speed  progressively;  the 
reverse  action  begins  at  zero  and  increases  in  speed.  Finally  the 
two  speeds  must  become  equal,  and  at  that  point  perceptible  change 
in  the  condition  of  the  whole  must  cease  (cf.  pp.  88-89). 

The  most  immediate  inference  from  this  mode  of  viewing  the 
matter  is,  that  the  apparent  halt  in  the  progress  of  the  action 
does  not  indicate  any  cessation  of  either  chemical  change.  Both 
changes  must  go  on,  in  consequence  of  the  continued  encounters  of 
the  proper  molecules.  But  since  the  two  changes  proceed  with 
equal  speeds  they  produce  no  alteration  in  the  mass  as  a  whole.  In 
fact,  the  final  state  is  one  of  equilibrium,  and  not  of  rest,  one  of 
balanced  activity  and  not  of  repose.  Hence,  chemical  changes 
which  are  reversible  lead  to  that  condition  of  seemingly  suspended 
action  which  we  speak  of  as  chemical  equilibrium. 

Chemical  Equilibrium  and  its  Characteristics.  —  The  de- 
tailed discussion  of  the  relations  of  liquid  and  vapor  (pp.  78,  87- 
90),  and  of  saturated  solution  and  undissolved  solid  (pp.  127,  130- 
133),  has  already  familiarized  us  with  the  term  equilibrium  and  its 
significance.  We  can,  in  fact,  apply  to  the  discussion  of  any  kind 
of  reversible  phenomena,  the  sets  of  ideas  in  regard  to  exchanges  of 
molecules  there  elaborated. 

In  particular,  the  reader  will  note  that  the  three  characteristics  of 
a  state  of  equilibrium,  developed  and  illustrated  in  the  case  of  the 
physical  equilibrium  between  a  liquid  and  its  vapor  (p.  89),  apply 
also  to  a  typical  case  of  chemical  equilibrium,  such  as  that  in 
Deacon's  process  now  before  us.  Thus: 


180  COLLEGE    CHEMISTRY 

1.  There   are   the  two  opposing   tendencies,   which  ultimately 
balance  one  another.     Here  they  are  the  tendency  of  the  steam  and 
chlorine  to  produce  hydrogen  choride  and  oxygen,  and  the  tendency 
of  the  hydrogen  chloride  and  oxygen  to  reproduce  steam  and 
chlorine  by  this  interaction.     In  other  words,  they  are  the  apparent 
activity  of  the  hydrogen  chloride  and  oxygen  interaction,  and  the 
apparent  activity  *  of  the  steam  and  chlorine  reaction. 

2.  At  equilibrium  the  two  opposing  tendencies  or  activities  are  still 
in  full  operation,  although  their  effects  then  neutralize  one  another. 

3  (and  this  is  the  chief  mark  of  chemical,  as  it  is  of  physical 
equilibrium).  The  system  is  in  a  sensitive  state,  so  that  a  change 
in  the  conditions  (temperature  and  pressure  or  concentration),  even 
if  slight,  produces  a  corresponding  change  in  the  state  of  the  system, 
and  does  this  by  favoring  or  disfavoring  one  of  the  two  opposing 
tendencies  or  apparent  activities.  Such  a  change  is  called  a  dis- 
placement of  the  equilibrium,  for  the  system  settles  down  in  a  new 
state  of  equilibrium  with  new  proportions  of  the  two  sets  of  sub- 
stances, corresponding  to  the  changed  conditions.  Thus,  in  the 
present  instance,  a  change  from  345°,  where  there  is  80  per  cent 
of  the  material  in  the  form  of  steam  and  chlorine,  to  384°  results 
in  the  diminution  of  this  proportion  to  75  per  cent.  The  equilib- 
rium is  affected  by  changes  in  concentration  also,  as  we  shall 
presently  see  (pp.  181,  186). 

Now,  the  foregoing  facts  show  that  the  key  to  understanding 
chemical  activities,  their  magnitudes,  their  changes,  and  especially 
their  practical  results,  must  lie  in  knowing  how  changes  in  the 
conditions  affect  them.  Hence,  to  the  chemist,  familiarity  with  the 
influence  of  conditions  on  chemical  phenomena  must  be  of  the 
greatest  practical  importance.  We  therefore  address  ourselves 
to  the  discussion  of  this  subject. 

The  "conditions"  to  be  considered  are  familiar,  —  temperature, 
and  concentration  or,  in  the  case  of  a  gas,  partial  pressure.  The 
"activity"  of  an  action  is  accurately  measured  by  the  speed  with 

*  We  use  the  term  "apparent  activity"  for  the  activity  as  we  see  it.  In 
the  same  action  it  varies  with  the  conditions.  The  intrinsic  activity  or  affinity, 
on  the  other  hand,  is  the  absolute  activity  of  the  action  irr&peclive  of  condi- 
tions. Its  value  can  be  determined  only  by  eliminating  the  effect  of  conditions, 
a  matter  which  is  too  abstract  for  consideration  here.  The  apparent  activity 
is  the  practical  thing  which  we  observe. 


CHEMICAL   EQUILIBRIUM  181 

which  the  action  proceeds.  Thus,  if  the  foregoing  section  be  re- 
examined,  it  will  be  seen  that  we  spoke  throughout  of  the  speed, 
rather  than  of  the  tendency  or  activity. 

Finally,  temperature  and  other  conditions  influence  also  the 
activities  in,  and  therefore  the  speeds  of,  those  actions  which  pro- 
ceed to  completion,  and  are  not  reversible.  Hence,  unless  our 
statements  are  expressly  restricted  to  reversible  actions  and  to 
states  of  equilibrium,  they  apply  to  all  chemical  changes. 

The  Influence  of  Concentration.  —  In  the  first  place,  let  us 
assume  that  the  temperature  is  constant,  and  let  us  confine  our 
attention  for  the  present  to  the  influence  of  concentration  upon  a 
chemical  reaction.  We  have  seen  (p.  178)  that  the  speed  of  a 
chemical  change  is  determined  by  the  frequency  with  which  the 
molecules  of  the  interacting  substances  encounter  one  another. 
The  frequency  of  the  encounters  amongst  a  given  set  of  molecules, 
resulting  in  a  definite  chemical  change,  will  in  turn  evidently 
depend  entirely  upon  the  degree  to  which  the  molecules  are  con- 
centrated in  each  other's  neighborhood.  Larger  amounts  of  one 
of  the  materials,  for  example,  will  not  result  in  more  rapid  chemical 
action,  if  the  larger  amount  of  material  is  also  scattered  through  a 
larger  space.  Chemical  changes,  therefore,  are  not  accelerated  by 
increasing  the  mere  quantity  of  any  ingredient,  but  only  by  in- 
creasing the  concentration  of  its  molecules.  Thus,  a  large  amount 
of  hydrochloric  acid  with  a  piece  of  zinc  will  generate  hydrogen  no 
faster  than  a  smaller  amount.  But  substitution  of  more  concen- 
trated acid  will  instantly  increase  the  speed  of  the  action.  In  the 
second  case,  the  number  of  molecules  of  the  acid  reaching  the  zinc 
per  second  is  greater,  and  this  action,  being  non-reversible,  pro- 
ceeds more  rapidly  to  complete  consumption  of  the  zinc.  So  also, 
iron  burns  faster  in  oxygen  (100  per  cent)  than  in  air  (20  per  cent 
oxygen). 

With  a  reversible  action  the  effect  on  the  speed  is  the  same,  ex- 
cepting that  the  continued  activity  of  the  reverse  action  prevents 
the  direct  one  from  reaching  completion. 

Thus,  if,  in  the  action  of  hydrogen  chloride  upon  oxygen,  we 
introduce  into  the  same  space  an  extra  amount  of  oxygen,  this 
facilitates  the  formation  of  steam  and  chlorine  by  increasing  the 
possibilities  of  encounter  between  molecules  of  hydrogen  chloride 


182  COLLEGE    CHEMISTRY 

and  oxygen.  At  the  same  time  it  does  not  affect  (cf.  p.  72)  the 
number  of  encounters  in  a  given  time  of  steam  and  chlorine  mole- 
cules with  one  another  which  result  in  the  reverse  transformation. 
The  proportion  of  chlorine  (and  steam)  formed,  therefore,  from  a 
given  amount  of  hydrogen  chloride  will  be  greater,  although  the 
total  possible  (by  complete  consumption  of  the  materials)  has  not 
been  altered,  since  the  quantity  of  one  ingredient  only  has  been 
increased.  The  introduction  of  an  excess  of  hydrogen  chloride 
would  have  had  precisely  the  same  effect. 

An  Experimental  Illustration.  —  A  reaction  in  which  the 
effects  of  different  concentrations  were  carefully  studied  by  Glad- 
stone (1855)  affords  a  good  illustration.  If  ferric  chloride  and 
ammonium  thiocyanate  are  mixed  in  aqueous  solution,  a  liquid 
containing  the  soluble,  blood-red  ferric  thiocyanate  is  produced. 
The  compound  radicals  are  (NBU)  and  (CNS),  and  the  action  is 
a  simple  double  decomposition: 

FeCU  +  3  NH4CNS  ?±  Fe(CNS)3  +  3  NH4C1. 

The  action  is  a  reversible  one,  and  the  mixture  is  homogeneous,  i.e., 
there  is  no  precipitation.  Now,  if  the  two  just-named  salts  are 
mixed  in  very  dilute  solution  in  the  proportions  required  by  the 
equation,  say  by  adding  20  c.c.  of  a  decinormal  solution  (p.  124) 
of  each  salt  to  several  liters  of  water,  a  pale-reddish  [solution  is 
obtained.  When  this  is  divided  into  four  parts,  and  one  is  kept  for 
reference,  the  addition  of  a  little  of  a  concentrated  solution  of  ferric 
chloride  to  one  jar,  and  of  ammonium  thiocyanate  to  another,  will 
be  found  to  deepen  the  color  by  producing  more  of  the  ferric  thio- 
cyanate. On  the  other  hand,  mixing  a  few  drops  of  concentrated 
ammonium  chloride  solution  with  the  fourth  portion  will  be  found 
to  remove  the  color  almost  entirely,  on  account  of  its  influence  in 
accelerating  the  backward  change. 

The  Law  of  Molecular  Concentration.  —  The  general  prin- 
ciple discussed  and  illustrated  in  this  section  may  be  called  the  law 
of  molecular  concentration,  and  may  be  stated  as  follows:  In 
every  chemical  change  the  apparent  activity,  and  therefore  the  speed 
of  the  action,  is  proportional  to  the  molecular  concentration  of  each 
interacting  substance.  This  holds  whether  the  action  is  reversible 
or  not. 


CHEMICAL   EQUILIBRIUM  183 

We  shall  next  give  a  more  precise,  semi-mathematical  formula- 
tion of  this  law,  as  this  formulation  will  be  of  use  later,*  and  then 
proceed  to  illustrate  the  application  of  the  law,  by  showing  how 
it  explains  large  classes  of  actions  of  which  we  have  already  en- 
countered many  examples. 

^Formulation  of  the  Law  of  Molecular  Concentration.  — 

The  mathematical  formulation  of  the  law  describing  the  influence 
of  the  concentration  of  the  molecules  of  each  participating  sub- 
stance upon  the  speed  of  the  action  is  extremely  simple.  When  the 
actual  concentrations  of  the  molecules  are  specified  (in  moles, 
pp.  102,  125,  per  liter),  and  the  speed  is  suitably  expressed  (in  moles 
transformed  per  minute  or  per  hour),  we  find  that  the  speed  is 
proportional  to  the  concentration  of  each  molecule  appearing  in 
the  molecular  equation  for  the  action.  Thus  in  the  interaction  of 
hydrochloric  and  hypochlorous  acids  (the  reverse  of  the  action  of 
chlorine  on  water,  pp.  161,  178),  if  [HC1]  and  [HC1O]  represent  the 
concentrations  of  the  molecules  HC1  and  HC10,  and  A;  is  a  con- 
stant, and  S  is  the  speed,  then 

[HC1]  X  [HC10]  X  k  =  8. 

t 

Again,  for  the  dissociation  of  phosphorus  pentachloride  vapor 
into  phosphorus  trichloride  and  chlorine  (p.  117):  PCU  — >  PCls  + 
C12,  if  [PC15]  represents  the  concentration  of  the  PC15  molecules,  &i 
is  a  constant,  and  Si  is  the  speed  of  decomposition: 

[PC15]  X  h  =  Si. 

Similarly,  for  the  reverse  action:  PC13  +  C12  ->  PC15,  if  [PCls]  and 
[C12]  stand  for  the  molecular  concentrations  of  these  substances: 

[PC13]  X  [C12]  X  h  =  82. 

The  constant  has  a  different  value  in  each  separate  action.  It 
includes  the  value  of  the  intrinsic  affinity  or  activity  of  the  sub- 
stances, and  the  catalytic  effect  (p.  29),  if  any,  of  the  materials 
present. 

*  This  formulation  of  the  law  is  not  required,  or  referred  to  in  the  sections 
which  follow.  The  section  and  the  following  one  may  therefore  be  omitted 
for  the  present  and  be  taken  up  in  connection  with  Chaps.  XX  and  XXXV. 


184  COLLEGE    CHEMISTRY 

Formulation  of  the  Condition  for  Chemical  Equilibrium. 

—  The  foregoing  plan  may  be  used  further  to  formulate  the  con- 
dition for  chemical  equilibrium.  As  we  have  seen  (p.  179),  a 
characteristic  of  a  system  in  chemical  equilibrium  is  that  the  speeds 
of  the  forward  and  reverse  actions  have  become  equal.  If,  then, 
[PCl5]eqm.,  [PCl3]eqm.,  and  [Cl2]eqm.  now  represent  the  molecular  con- 
centrations when  the  system  has  reached  equilibrium,  then,  since 
the  speeds  are  equal: 

[PCl3]eqm.  X   [Cl2]eqm.  X  fe  =   [PClJeqm.  X  fe, 
[PClsleqm.  X  [Cl2]eqm.  _  fel  _  mnt,tflnf 

"      " 


[PCl6] 


eqm. 


In  words,  this  means  that  if  we  change  the  amount  of  the  penta- 
chloride  placed  in  the  vessel,  or  if  we  use  amounts  of  chlorine  and 
trichloride  which  are  not  equivalent,  the  numerical  value  at  equilib- 
rium of  each  concentration  ([PCls]  etc.)  will,  of  course,  be  differ- 
ent, but  the  product  of  the  concentrations  of  trichloride  and 
chlorine,  divided  by  the  concentration  of  the  pentachloride,  will 
always  give  the  same  numerical  value  for  the  constant  at  the  same 
temperature.  This  numerical  value  is  called  the  equilibrium 
constant. 

Applications:  The  Forward  Action.  Homogeneous  and 
Inhomogeneous  Systems.  —  While  there  are  all  degrees  of  speed 
in  chemical  actions,  yet  in  practice  we  quickly  distinguish  two 
different  classes.  There  is  a  class  of  actions  of  which  most  exam- 
ples are  almost  instantaneously  accomplished,  and  a  class  in  which, 
frequently,  the  operation  takes  minutes  or  even  hours.  The 
classes  overlap,  but,  in  a  general  way,  the  following  distinction  may 
be  made. 

To  the  former,  speedy  class  belong  the  explosion  of  hydrogen 
and  oxygen  or  other  gaseous  mixtures,  and  the  interactions  when 
solutions  are  mixed,  as  in  precipitations.  In  view  of  the  foregoing 
explanations,  we  perceive  that  the  rapid  accomplishment  of  such 
actions  is  due,  not  so  much  to  any  especially  great  intrinsic  affinity, 
as  to  the  homogeneous  state  of  mixture  of  the  interacting  materials. 
This,  of  course,  is  a  purely  physical,  and  not  a  chemical  motive  for 
speedy  interaction.  In  intimate  mixtures,  every  molecule  has  an 
equal  opportunity  freely  to  encounter  every  other  molecule  and 


CHEMICAL   EQUILIBRIUM  185 

there  is  therefore  no  mechanical  impediment  to  the  operation  of 
the  affinities  of  the  substances.  Hence  the  apparent  activity  is 
great. 

To  the  second  class,  comprising  the  slower  actions,  belong  cases 
like  the  interaction  of  a  piece  of  zinc  with  hydrochloric  acid,  or  of 
manganese  dioxide  (p.  158)  with  the  same  acid,  whereby  hydrogen 
and  chlorine,  respectively,  are  slowly  evolved,  and  the  solid  is  grad- 
ually consumed.  Here  the  hindrance  is  evidently  the  fact  that 
the  interacting  substances  are  not  intimately  mixed.  In  the  slow 
actions,  the  system  is  inhomogeneous.  Pulverizing  the  solid  before 
use  will  increase  the  speed,  indeed,  by  providing  more  surface  and 
better  mutual  contact,  but  will  not  transfer  the  action  to  the  rapid 
class.  It  is  chiefly  the  dissolved  part  of  the  substance  which  inter- 
acts, for  chemical  action  takes  place  between  molecules,  and  only 
the  dissolved  part  is  disintegrated  in  such  a  way  that  the  molecules 
are  readily  accessible.  Thus,  the  action  is  held  back  by  continual 
waiting  for  the  slow  replenishment,  from  the  "insoluble"  solid,  of 
the  supply  of  dissolved  molecules.  In  the  cases  cited,  the  restrain- 
ing influence  of  the  dissolving  process,  which  is  part  of  the  whole 
phenomenon,  may  be  formulated  thus: 

Zn(solid)  ±=>  Zn(dslvd)  +  2HC1  -» ZnCl2  +  H2. 
MnO2(solid)  ±^  MnO2(dslvd)  +  4HC1  -»  MnCl2  +  2H20  +  C12. 

Here,  again,  the  mechanical  details,  depending  on  physical  prop- 
erties, have  more  to  do  with  the  progress  of  the  action  than  has 
the  chemical  affinity.  In  terms  of  the  law  of  concentration,  the 
action  is  slow,  and  the  apparent  activity  small,  because  the  con- 
centration of  the  acting  molecules  of  one  of  the  substances  is  very 
small,  and  cannot  be  increased  because  of  low  solubility. 

Applications:  The  Reverse  Action.  Displacement  of 
Equilibria.  —  We  have  seen  (p.  182)  that  one  way  in  which  a  re- 
versible action  may  be  forced  nearer  to  completion,  in  one  direction 
or  the  other,  is  the  introduction  of  an  excess  of  one  of  the  ingre- 
dients contributing  to  the  forward  action.  This  method  of  dis- 
placing the  equilibrium  point,  however,  cannot  be  very  effective, 
unless  it  is  possible  to  introduce  an  exceedingly  large  excess  of  the 
selected  ingredient  in  a  high  degree  of  molecular  concentration, 
since  this  operation  does  not  in  any  way  affect  or,  in  particular, 


186  COLLEGE    CHEMISTRY 

restrain  the  reverse  action  which  is  continually  undoing  the  work  of 
the  forward  one.  A  much  more  effective  means  of  furthering  the 
desired  direction  of  such  actions  is  found,  therefore,  in  the  restraint 
or  practical  annulment  of  the  reverse  action.  A  good  way  of 
accomplishing  this  is  to  allow  the  products  of  the  direct  action  to 
separate  into  an  inhomogeneous  mixture.  Any  agency  which  could 
remove  the  water  vapor  as  fast  as  it  was  formed  by  the  interac- 
tion of  hydrogen  chloride  and  oxygen,  for  example,  would  entirely 
stop  the  reproduction  of  these  substances,  and  so  would  enable 
the  forward  action  (4HC1  +  O2  — >  2H20  +  C12)  to  run  to  completion. 
This  might  be  realized  by  causing  one  end  of  a  sealed  tube 
charged  with  the  substances,  after  the  contents  had  settled  down 

to  a  condition  of  equilibrium,  to 
project  from  the  bath  in  which  the 
whole  had  been  kept  at  345°  (Fig. 
74,  which  is  simply  diagrammatic). 
By  cooling  this  end,  a  large  part 
of  the  steam  would  quickly  be  con- 
densed in  it  to  the  liquid  form, 

while  the  other  substances  would  remain  gaseous.  In  other  words, 
the  concentration  of  the  water  vapor  would  be  greatly  reduced.  In 
fact,  only  the  trace  of  vapor  which  cold  water  gives  would  then  be 
available  to  interact  with  the  chlorine,  and  reproduce  hydrogen  chlo- 
ride. Meanwhile  the  decomposition  of  the  latter  would  go  on,  and 
thus,  eventually,  almost  all  the  water  would  be  found  in  one  end  of 
the  tube,  and  the  chlorine,  all  free,  would  occupy  the  rest.  By  this 
purely  mechanical  adjustment  the  chemical  change  would  therefore 
be  carried  from  80  per  cent  completion  to  almost  absolute  completion: 

4HC1  +  02  *±  2C12  +  2H20  (vapor)  *=»  2H20  (liq.) 
If,  on  the  other  hand,  arrangements  were  made  to  have  pow- 
dered marble,  in  a  sealed  bulb  of  thin  glass,  enclosed  in  the  tube, 
we  might  imagine  the  very  opposite  of  the  above  effect  to  be  pro- 
duced.    The  breaking  of  the  bulb  of  marble,  when  equilibrium 
had  been  reached,  would  provide  means  for  the  removal  of  all  the 
hydrogen  chloride,*  while  the  other  three  substances  would  still  be 
*  The  hydrogen  chloride  would  be  destroyed  by  interaction  with  the  marble: 

2HC1  +  CaC03  -» CaCl2  +  CO2  +  H2O. 

The  calcium  chloride  is  a  solid.  The  gas,  carbon  dioxide,  does  not  interact 
with  the  other  substances,  and  would  not,  therefore,  interfere  with  the  forma- 
tion of  fresh  hydrogen  chloride. 


CHEMICAL   EQUILIBRIUM  187 

gaseous.  Thus,  the  compound  (HC1)  having  been  reduced  in 
concentration  to  the  point  of  being  removed  entirely,  there  would 
be  no  direct  action  to  undo  the  work  of  the  reverse  action.  The 
whole  chlorine  would,  therefore,  soon  have  passed  through  the 
form  HC1.  Hence,  by  another  mechanical  arrangement,  an  action 
which  ordinarily  could  progress  to  only  20  per  cent  would  be 
turned  into  a  complete  one : 

2C12  +  2H20  <=»  2H20  +  4HC1  (+CaC03  -+  CaCl2  +  H20  +  C02). 

Reversibility  Usually  Avoided.  —  In  every-day  chemical 
work,  since  our  object  is  usually  to  prepare  some  one  substance, 
chemists  either  avoid  chemical  changes  which  are  notably  re- 
versible, or  adjust  the  conditions,  as  is  done  in  the  foregoing 
illustrations,  so  that  the  reverse  of  the  action  which  they  desire  is 
prevented.  In  consequence  of  this,  when  carrying  out  the  direc- 
tions for  making  familiar  preparations,  the  fact  that  such  actions 
are  reversible  at  all  very  readily  escapes  our  notice.  Arranging 
the  conditions  so  that  the  separation  of  a  solid  body  by  precipita- 
tion, or  the  liberation  of  a  gas,  takes  place,  are  the  two  commonest 
ways  of  rendering  a  reversible  action  complete.  Excellent  ex- 
amples of  both  of  these  are  furnished  by  the  chemical  change 
used  in  producing  hydrogen  chloride  by  the  interaction  of  salt  and 
sulphuric  acid,  the  full  discussion  of  which  (p.  142)  should  now  be 
studied  attentively  in  the  light  of  these  explanations. 

History.  —  The  conceptions  discussed  in  this  chapter  are  not 
new,  although  they  have  come  into  general  use  rather  recently. 
The  law  of  reaction  speed,  and  the  influence  of  the  concentrations 
of  the  reacting  substance  thereon  (p.  183),  was  set  forth  and 
formulated  by  Wilhelmy  as  early  as  1850.  Gladstone  (1855) 
studied  quantitatively  the  influence  of  concentration  in  cases  of 
chemical  equilibrium  (p.  182).  The  kinetic  explanation  (p.  178) 
was  developed  by  Williamson  (1851) .  Finally,  the  laws  of  chemical 
equilibrium  were  formulated  more  explicitly  and  applied  more 
thoroughly  by  two  Swedish  chemists,  Guldberg  and  Waage 
(1864-9). 

The  Influence  of  Temperature  on  the  Speed  of  any  Re- 
action. —  The  activity  of  chemical  change,  and  therefore  the 


188  COLLEGE   CHEMISTRY 

speed  of  all  chemical  changes,  is  increased  by  raising  the  temperature 
and  diminished  by  lowering  it  (cf.  p.  59).  Thus,  zinc  displaces 
hydrogen  more  rapidly  from  hot  than  from  cold  hydrochloric  acid. 
Different  actions  are  affected  in  different  degrees,  and  no  simple 
rule  accurately  denning  the  effect  can  be  given.  Roughly  speak- 
ing, however,  a  rise  of  10°  doubles  the  speed  of  every  action.  A 
rise  of  100°  will  therefore  make  the  speed  roughly  1024  times 
greater.  Hence,  when  the  chemist  finds  that  two  substances 
show  no  evidence  of  interaction,  he  infers  that  there  must  be  either 
slow  action  or  none,  and  he  seeks  to  settle  the  question  quickly  by 
heating  the  mixture. 

The  Influence  of  Temperature  on  a  System  in  Equilib- 
rium. —  In  a  reversible  change  the  two  opposing  reactions  are 
different  actions  and  their  speeds  are  therefore  affected  in  different 
degrees  by  the  same  alteration  in  temperature.  Hence,  when  the 
temperature  is  changed,  the  relative  amount  of  the  two  sets  of 
materials  present  is  altered  and  the  equilibrium  is  displaced.  Thus, 
in  Deacon's  process,  a  rise  of  40°  in  the  temperature  displaces  the 
equilibrium  backwards. (p.  180),  and  diminishes  the  yield  of  chlo- 
rine by  5  per  cent.  In  the  vapor  of  phosphorus  pentachloride 
(p.  117),  the  displacement  is  in  the  opposite  direction.  The  vapor 
is  a  mixture  of  the  pentachloride  with  the  trichloride  and  free 
chlorine:  PC18  <=±  PC13  +  C12.  At  200°,  51.5  per  cent  of  the 
material  is  present  as  pentachloride  and  48.5  per  cent  as  trichloride 
and  chlorine.  Raising  the  temperature  to  250°  changes  the  pro- 
portions to  20  per  cent  and  80  per  cent,  respectively.  At  300°  only 
3  per  cent  of  the  pentachloride  remains.  Evidently,  here,  raising 
the  temperature  favors  the  decomposition  of  the  pentachloride, 
and  therefore  increases  the  speed  of  its  dissociation  more  than  it 
does  the  speed  of  the  reunion  of  the  trichloride  and  chlorine. 

Van't  Hoff's  Law.  —  One  use  of  a  law  is  to  enable  us  to  answer 
a  question,  when  we  have  not  in  memory  the  fact  constituting  the 
answer,  and  even  when  we  have  never  read  or  heard  the  fact.  The 
law  or  rule  enables  a  little  reasoning  to  take  the  place  of  a  vast 
amount  of  memorizing.  Thus,  to  answer  the  question:  Does 
sodium  chloride  always  have  the  same  composition,  it  is  not 
necessary  to  have  read  and  to  remember  all,  or  any  of  the  numerous 


CHEMICAL   EQUILIBRIUM  189 

investigations  of  this  substance  that  have  been  made.  We  simply 
refer  the  question,  mentally,  to  the  law  of  definite  proportions, 
and  say  "yes."  Now  the  facts  mentioned  above  are  connected 
by  a  law  which  will  answer  many  practical  questions  in  chemistry. 
When  phosphorus  trichloride  and  chlorine  combine  (to  form 
PC15),  heat  is  given  out.  Conversely,  when  phosphorus  penta- 
chloride  dissociates,  heat  is  absorbed: 

PC15  +  30,000  cal.  <F±  PC13  +  C12. 

Now,  when  the  temperature  is  raised,  the  action  proceeds  in  the 
direction  of  decomposing  more  of  the  pentachloride.  That  is,  the 
equilibrium  is  displaced  in  the  direction  which  absorbs  heat. 

In  Deacon's  process,  we  find  that  the  interaction  of  hydrogen 
chloride  and  oxygen  liberates  heat, 

4HC1  +  02  <±  2H2O  +  2C12  +  28,000  cal., 

and  in  this  action  raising  the  temperature  drives  the  equilibrium 
backwards,  and  a  lowering  in  the  temperature  is  required  to  increase 
the  yield  of  chlorine. 

The  rule  is  obvious,  and  applies  to  all  reversible  reactions: 
When  the  temperature  of  a  system  in  equilibrium  is  raised,  the  equi- 
librium point  is  displaced  in  the  direction  which  absorbs  heat.  In 
other  words,  a  rise  in  temperature  favors  the  interaction  of  that 
one  of  the  two  sets  of  materials  to  which  the  heat  is  added  (+  sign) 
in  the  equation.  If  the  equation  happens  to  be  written  with  a 
negative  heat  of  reaction  (e.g.,  p.  175),  the  heat  can,  of  course,  be 
transferred  to  the  other  side  with  its  sign  changed.  This  law  is 
known  as  Van't  Hoff's  law  of  mobile  equilibrium. 

This  law  is  of  practical  value.  More  than  once,  in  chemical 
factories,  much  time  and  money  have  been  spent  on  trying  to 
arrange  machinery  to  give  a  better  yield  of  some  substance  at  a 
high  temperature,  when  a  reference  to  this  law  would  have  shown 
that  the  chief  change  necessary  was  to  use  a  lower  temperature. 
We  shall  frequently  have  occasion  to  refer  to  this  law. 

Application  to  Physical  Equilibria.  —  Van't  Hoff's  law 
applies  also  to  physical  processes.  Thus,  as  the  temperature  rises, 
a  substance  which  absorbs  heat  in  dissolving  will  become  more 
soluble.  This  is  the  commoner  case,  as  is  shown  by  the  way  in 


190  COLLEGE    CHEMISTRY 

which  most  solubility  curves  (Fig.  58,  p.  131)  ascend  with  rising 
temperature.  Conversely,  a  substance  which  gives  out  heat  in 
dissolving  is  less  soluble  with  rising  temperature  in  a  solution 
already  almost  saturated  with  the  compound.  For  example,  anhy- 
drous sodium  sulphate  gives  out  heat  in  dissolving,  and  so  its 
solubility  diminishes  (Fig.  59,  p.  132),  with  rising  temperature. 

Again,  the  vaporization  of  a  liquid  absorbs  heat,  and  so  an  in- 
crease in  temperature  will  increase  the  pressure,  and  therefore  the 
concentration  of  its  vapor  (p.  87). 

Le  Chatelier's  Law.  —  The  above  mentioned  law  is  really  a 
particular  case  of  a  more  general  one.  If  some  stress  (e.g.,  by 
change  of  temperature,  pressure,  or  concentration)  is  brought  to 
bear  on  a  system  in  equilibrium,  the  equilibrium  is  displaced  in  the 
direction  which  tends  to  undo  the  effect  of  the  stress.  Thus,  raising 
the  temperature  furthers  the  change  which  absorbs  heat  —  and 
therefore  would  tend  to  lower  the  temperature.  Increasing  the 
concentration  of  the  molecules  pushes  the  action  in  the  direction 
which  uses  up  these  very  molecules  (p.  181).  Again  pressure 
causes  ice  to  melt,  because  the  water  which  is  formed  occupies  a 
smaller  volume,  and  this  change  tends  to  relieve  the  pressure. 
But  pressure  will  not  cause  most  substances  to  melt,  because 
usually  the  liquid  form  occupies  a  greater  volume  and  its  produc- 
tion would  tend  to  increase  pressure. 

Summary.  —  In  this  chapter  we  have  answered  three  ques- 
tions: 

1.  Why  do  some  chemical  actions  cease,  while  still  incomplete? 
Answer:  They  are  reversible. 

2.  What  explains  the  position  of  the  equilibrium  point?    An- 
swers: (a)  Equal  effects  of  opposed  molecular  actions;  (6)  Equality 
in  speed  of  opposed  reactions. 

3.  What  will  displace  the  equilibrium  point?    Answer:    (a) 
Change  in  concentration  of  one  (or  more)  of  the  substances;   (6) 
Change  in  the  temperature. 

Exercises.  —  1.  Explain  the  completeness  of  the  action  by 
which  hydrogen  chloride  and  water,  respectively,  are  formed  by 
direct  union  of  the  elements. 


CHEMICAL   EQUILIBRIUM  191 

2.  Explain  the  completeness  of  the  action  by  which  silver 
chloride  (p.  148)  is  formed. 

3.  Explain  why  the  decomposition  of  potassium  chlorate  is 
complete. 

4.  In  view  of  the  statement  on  p.  20,  explain  why  mercuric  oxide 
is  completely  decomposed  by  heating.     Point  out  the  resemblance 
between  this  experiment  and  the  imaginary  one  illustrated  in 
Fig.  74  (p.  186). 

5.  Why  can  magnetic  oxide  of  iron  be  reduced  completely  by 
a  stream  of  hydrogen  (p.  59),  and  iron  oxidized  completely  by  a 
current  of  steam  (p.  51)? 

6.  With  the  phosphorus  pentachloride  system,  say  at  250°,  what 
effect  would  suddenly  enlarging  the  space  containing  a  given 
amount  of  the  vapor  produce?     What  would  be  the  effect  of  di- 
minishing the  space?    What  would  be  the  effect  of  introducing 
additional  chlorine  into  the  same  space  (p.  181)? 

7.  By  what  practical  means  could  the  degree  of  dissociation  of 
sulphur  vapor  (S8)  be  reduced,  without  changing  the  temperature 
(p.  117)? 

8.  What  inference  should  you  draw  from  the  fact  that:   (a)  the 
solubilities  of  potassium  nitrate  and  of  Glauber's  salt  (p.  132) 
increase  with  rise  in  temperature;    (6)  that  those  of  calcium  hy- 
droxide (p.  130)  and  triethylamine  decrease  with  rise  in  tempera- 
ture? 


CHAPTER  XV 
THE  HALOGEN   FAMILY 

THE  elements  to  which  we  have  so  far  devoted  most  attention 
have  been  oxygen,  hydrogen,  and  chlorine.  If  we  recall  the  chemi- 
cal properties  and  relations  of  these  elements  we  shall  recognize 
the  fact  that  they  all  possess  very  distinct  individualities. 

The  Chemical  Relations  of  Elements.  —  Hydrogen  is  a  sub- 
stance (p.  58)  which  unites  readily  with  oxygen  and  chlorine, 
less  readily  with  other  non-metals,  and  scarcely  at  all  with  metals. 
Oxygen  and  chlorine  resemble  each  other  somewhat  in  the  great- 
ness of  then-  chemical  activity  and  the  variety  of  free  elements  with 
which  they  are  capable  of  uniting,  but  differ  markedly  in  what  we 
have  called  their  chemical  relations  (p.  163).  The  resulting  com- 
pounds belong,  in  fact,  to  quite  different  classes  —  oxygen  forms 
oxides,  chlorine  forms  chlorides  —  and  elements  are  considered 
similar  only  when  they  resemble  each  other  in  chemical  relations, 
and  produce,  by  combination  with  the  same  element,  compounds 
having  similar  chemical  properties.  Thus,  the  common  oxide  of 
hydrogen,  water,  is  a  neutral  substance,  and  is  chemically  rather  in- 
different. The  chloride  of  hydrogen  in  aqueous  solution  is  a  strong 
acid  and  is  chemically  very  active.*  If  all  the  other  chemical  ele- 
ments differed  from  one  another  as  much  as  do  these  three,  they 
would  be  incapable  of  classification.  In  reality,  however,  we  find 
that  the  elements  can  be  grouped  together  in  sets.  They  are  classi- 
fied according  to  the  kind  of  substances  with  which  they  combine 
and  the  chemical  nature  of  the  products.  In  some  families  the  re- 
semblance is  close,  in  others  less  close.  The  present  group  is  of 
the  former  class,  and  will  serve,  therefore,  as  a  convenient  begin- 

*  The  difference  between  oxides  and  chlorides  is  seen  in  their  behavior. 
Thus,  oxides  often  unite  with  water  to  form  acids  or  bases  (p.  94).  Chlorides 
do  not  unite  with  water  to  form  new  substances  with  marked  characteristics 
(cf.  p.  96). 

192 


BEOMINE  193 

ning  in  the  work  of  tracing  relations  between  the  elements  and  in 
classifying  the  facts  of  descriptive  chemistry. 

The  Chemical  Relations  of  the  Halogens.  —  The  bromide 
(NaBr),  iodide  (Nal),  and,  to  a  less  extent,  the  fluoride  (NaF)  of 
sodium,  resemble  sodium  chloride  (NaCl)  in  appearance  and  be- 
havior. From  this  fact,  chlorine,  bromine,  iodine,  and  fluorine  are 
known  as  the  halogens  (Gk.,  salt  producers),  and  their  compounds 
are  named  the  halides.  The  halogens,  as  the  above  formulae  show, 
are  univalent.  They  all  form  compounds  with  hydrogen,  and 
these  compounds  closely  resemble  hydrogen  chloride.  For  ex- 
ample,  they  are  colorless,  they  are  gases  (except  hydrogen  fluoride, 
a  very  volatile  liquid),  they  are  very  soluble  in  water,  and  their 
solutions  are  acids.  Other  relations  will  be  given  in  a  summary 
at  the  end  of  the  chapter. 

BROMINE  Br2 

Occurrence.  —  The  compounds  of  chlorine,  bromine,  and 
iodine  usually  occur  together  in  nature,  while  the  compounds  of 
fluorine  are  not  found  in  the  same  sources.  Bromine  occurs 
chiefly  in  the  form  of  the  bromides  of  sodium  and  magnesium,  in 
the  upper  layers  of  the  natural  beds  of  rock  salt.  Liebig  made  it 
from  this  source  and  a  little  later  Bal- 
lard  (1826)  made  it  also  and  recognized 
it  as  a  new  element. 


Preparation.  —  In    the    chemical 
point  of  view  there  are  three  distinct 

ways  in  which  bromine  is  made.     1.  The  first  of  these  is  closely 
related  to  the  common  method  of  preparing  chlorine  (p.   158). 
As  hydrobromic  acid,  unlike  hydrochloric  acid,  is  not  formed  ex- 
tensively in  connection  with  any  chemical  industry,  potassium 
bromide  is  treated  in  a  retort  (Fig.  75)  with  concentrated 
phuric  acid,  and  the  product  is  oxidized  with  powdered  manganes 
dioxide   in   one    operation.      (For  equation   see   next  section.) 
Bromine  being  a  volatile  liquid,  while  the  sulphates  of  potassium 
and  manganese  are  involatile,  its  vapor  passes  off  when  the  s 
mixture  is  heated.     It  is  condensed  in  a  flask  surrounded  by  cold 


water. 


194  COLLEGE    CHEMISTRY 

2.  The  second  method  of  preparing  bromine  depends  on  the  fact 
that  chlorine  is  a  more  active  element  and  displaces  bromine  from 
combination.     When,  therefore,  chlorine  is  passed  into  a  solution 
of  potassium  or  sodium  bromide,  potassium  or  sodium  chloride  is 
formed  and  the  bromine  liberated: 

2NaBr  +  C12  ->  2NaCl  +  Br2. 

When  the  liquid  is  warmed,  the  bromine  passes  off  along  with  a 
part  of  the  water,  and  may  be  condensed  as  before. 

3.  Aqueous  solutions  of  soluble  bromides  may  be  decomposed 
by  means  of  a  current  of  electricity.     The  bromine  is  set  free  at  the 
positive  electrode. 

Commercial  Extraction.  —  Two-thirds  of  the  world's  supply 
is  obtained  from  Stassfurt,  where,  after  the  extraction  of  the 
potassium  chloride  from  the  impure  carnallite  (KCl,MgCl2,6H20), 
the  mother-liquor  is  found  to  contain  the  more  soluble  sodium 
and  magnesium  bromides  in  considerable  quantities.  The  warm 
mother-liquor  trickles  down  over  round  stones  in  a  tower.  The 
chlorine  is  introduced  from  below  and  dissolves  in  the  liquid.  The 
bromine  is  thus  liberated  and  passes  off  as  vapor.  A  part  of  our 
supply  of  bromine  is  obtained  from  the  brines  of  Ohio,  West 
Virginia,  and  Kentucky,  from  which,  after  most  of  the  common 
salt  has  been  removed  by  crystallization,  the  bromine  is  obtained 
by  the  first  method.  In  Michigan  the  brines  are  treated  with 
electrolytic  chlorine  by  the  second  method. 

Partial  Equations,  a  Plan  for  Making  Complex  Equations. 

—  When  an  equation  involves  more  than  two  initial  substances  or 
products,  as  does  the  one  for  the  first  method  of  preparing  bromine, 
it  cannot  readily  be  worked  out  by  the  method  formerly  recom- 
mended (p.  51).  After  the  formulae  of  all  the  substances,  on  both 
sides,  have  been  set  down,  it  is  difficult  to  hit  upor>  the  proper  co- 
efficients required  to  balance  the  equation.  In  such  cases,  a  good 
plan  is  to  select  two  of  the  initial  substances,  and  make  a  partial 
equation  showing  part  of  the  action  and  including  at  least  one 
actual  product.  Any  unused  units  (not  constituting  a  product) 
are  then  set  down  also  and  treated  as  a  balance.  Thus  the  first 


BROMINE  195 

two  of  the  substances  named  will  furnish  potassium-hydrogen 
sulphate : 

Partial,  1 :          KBr  +  H2S04  -»  KHS04  (+  HBr) .  (1) 

Similarly,  the  manganese  dioxide  and  sulphuric  acid  will  give 
manganous  sulphate: 

Partial,  2:     MnO2  +  H2S04  -»  MnS04  +  H20  (+  0).  (2) 

We  then  perceive  that  the  bromine  must  come  from  the  oxidation 
of  the  first  balance  (HBr)  by  the  second  (0) : 

Partial,  3:  (2HBr)  +  (O)  ->  H20  +  Br2.  (3) 

The  third  partial  equation  shows  that  2HBr  will  be  needed  for  the 
amount  of  O  obtainable  from  Mn02,  so  we  go  back  to  (1)  and 
multiply  it  by  two  throughout : 

2KBr  +  2H2S04  -»  2KHS04  (+  2HBr).  (1) 

MnO2  +  H2SO4  ->  MnS04  +  H20  (+  0).  (2) 

(2HBr)  +  (0)  ->  H20  +  Br2. (3) 

2KBr  +  3H2SO4  +  Mn02  ->  2KHS04  +  MnS04  +  2H20  +  Br2. 

When  we  now  add  the  real  substances  used  and  produced,  as  they 
occur  in  these  partial  equations,  and  leave  out  the  balances,  which 
have  been  adjusted  so  as  to  cancel  one  another,  we  obtain  the  final 
equation  for  the  action.  It  must  be  observed  that  this  subdivi- 
sion of  the  action  into  parts  is  a  purely  arithmetical  device.  It 
is  still  true,  however,  that  we  are  aided  in  the  selection  of  partial 
actions  at  each  step  by  following  some  plausible  theory  as  to 
stages  for  the  action  which  would  be  chemically  conceivable. 

Physical  Properties.  —  Bromine  is  a  dark-red  liquid  (sp.  gr. 
3.18).  It  boils  at  59°,  forming  a  deep-red  vapor,  and  even  at  ordi- 
nary temperatures  gives  a  high  vapor  pressure  (150  mm.  at  18°)  and 
evaporates  quickly.  When  cooled  it  forms  red,  needle-shaped 
crystals  (m.-p.  —  7.3°).  A  saturated  aqueous  solution  (bromine- 
water)  contains  3  parts  of  bromine  hi  100  parts  of  water.  The 
element  is  much  more  soluble  in  carbon  disulphide,  alcohol,  and 
other  organic  solvents.  Up  to  750°,  the  G.M.V.  weighs  160  g. 
(corresponding  to  Br2),  against  28.955  g.  for  air. 

Bromine  (Gk.,  a  stench)  has  a  most  pungent  odor.     It  has  a 


196  COLLEGE   CHEMISTRY 

very  irritating  effect  on  the  mucous  membrane  of  the  nostrils  and 
throat.    If  spilled  upon  the  hands  it  destroys  the  tissues  and  leaves 
sores  which  are  liable  to  infection. 
Free  bromine  has  no  effect  upon  starch  emulsion  (see  Iodine). 

Chemical  Properties.  —  A  jet  of  hydrogen  gas  burns  in 
bromine  vapor.  The  union  is  much  slower  than  in  the  case  of 
chlorine  (Heat  of  formation,  +  12,300  cal.). 

Bromine  forms  compounds  directly,  both  with  non-metals,  like 
phosphorus  and  arsenic,  and  with  most  of  the  metals,  which  catch 
fire  when  thrown  into  the  vapor.  In  all  cases  the  interaction  is 
less  violent  than  when  chlorine  is  used,  and  bromine  is  displaced 
from  combination  with  hydrogen  and  with  the  metals  by  free 
chlorine. 

Silver  bromide  is  the  sensitive  material  in  photographic  plates, 
and  potassium  and  sodium  bromides  are  used  as  sedatives  in 
medicine.  Bromine  is  employed  in  the  preparation  of  organic 
dyes. 

HYDROGEN  BROMIDE  HBr 

Preparation.  —  It  might  be  expected  that  the  most  convenient 
way  of  producing  this  compound  would  be  similar  to  that  used  in 
preparing  hydrogen  chloride,  namely,  by  the  action  of  concentrated 
sulphuric  acid  upon  some  common  bromide,  such  as  potassium 
bromide  (KBr  +  H2SO4  +±  HBr  +  KHSO4).  Hydrogen  bromide 
being  less  stable,  however,  a  large  part  of  it  is  oxidized  by  the 
sulphuric  acid  and  the  product  is  mixed  with  sulphur  dioxide  and 
free  bromine. 

H2S04  +  2HBr  -*  2H20  +  S02 1  +  Br2 1 . 

Since  all  acids  decompose  all  salts  more  or  less,  use  of  an  acid 
which  does  not  give  up  its  oxygen  so  readily,  such  as  phosphoric 
acid,  will  yield  pure  hydrogen  bromide  (KBr  +  H3P04  — >  HBr  j 
+  KH2P04).  The  small  solubility  of  the  salt  in  concentrated 
phosphoric  acid  retards  the  interaction  and  makes  the  evolution  of 
the  gas  very  slow,  however. 

Pure  hydrogen  bromide  is  best  prepared  by  the  action  of  water 
upon  phosphorus  tribromide  (see  Hydrolysis,  below).  When 
bromine  and  phosphorus  are  mixed,  a  violent  union  of  the  two 


HYDROGEN    BROMIDE 


197 


elements  takes  place,  producing  phosphorus  tribromide  PBr3. 
This  substance,  which  is  a  colorless  liquid,  is  in  turn  broken  up 
with  great  ease  by  water,  producing  phosphorous  acid,  which  is  not 
volatile,  and  gaseous  hydrogen  bromide: 


OH 


OH 


OH 


OH 


FIG.  76. 


In  practice,  these  two  actions  are  carried  on  simultaneously.  To 
diminish  the  vigor  of  the  interaction,  red  phosphorus  is  taken  in- 
stead of  yellow,  and  is  mixed  with  two  or  three  times  its  weight  of 
sand  in  a  flask  (Fig.  76) .  A  small  quantity  of  water  is  added.  Ex- 
cess of  water  must  be  avoided,  as  the  hydrogen  bromide  produced 
is  extremely  soluble,  and  would  there- 
fore be  retained  in  the  flask  instead  of 
being  disengaged  as  gas.  The  bromine 
is  placed  in  the  dropping  funnel,  and 
admitted,  a  little  at  a  time,  to  the 
mixture.  The  gas  produced  is  passed 
through  a  U-tube  containing  red  phos- 
phorus mixed  with  glass  beads.  The 
phosphorus  combines  with  any  free 
bromine  carried  along  with  the  gas. 
The  second  U-tube,  containing  water,  may  be  attached  when  a 
solution  of  the  gas  is  required.  The  gas  may  be  collected  in  a 
jar  by  upward  displacement  of  air. 

Hydrolysis.  —  The  interaction  of  water  with  phosphorus  tri- 
bromide (foregoing  section)  illustrates  an  important  property  of 
water  (p.  92) .  The  action  is  a  double  decomposition  in  which  water 
is  one  of  the  interacting  substances  and  is  called  an  hydrolysis  (Gk., 
loosening  by  water).  The  water  divides  into  the  radicals  H  and 
OH,  and  the  former  unites  with  the  more  active  non-metal  in  the 
substance  (the  bromine,  in  PBr3)  and  the  hydroxyl  with  the  other 
element.  For  example,  PC13  +  3HOH  -»  P(OH),  +  3HC1.  All 
the  halides  of  the  non-metals  are  thus  hydrolyzed,  as  area 
other  classes  of  compounds. 

Physical  Properties.  — Hydrogen  bromide  is  a  colorless  gas 
with  a  sharp  odor.     It  is  two  and  a  half  times  as  heavy  as  air. 


198  COLLEGE   CHEMISTRY 

easily  reduced  to  the  liquid  condition  (b.-p.  —  69°).  It  is  ex- 
ceedingly soluble  in  water,  and  in  contact  with  moist  air  condenses 
the  water  vapor  to  clouds  of  liquid  particles.  Pure  hydrogen 
bromide,  whether  in  the  gaseous  condition  or  in  the  liquefied  form, 
is  a  nonconductor  of  electricity  (see  below). 

Chemical  Properties.  —  The  properties  are  like  those  of 
hydrogen  chloride  (p.  145).  It  is  somewhat  less  stable,  and  dis- 
sociation begins  to  be  noticeable  at  800°.  When  free  from  water, 
it  is  not  an  acid  (see  below).  The  gas  interacts  vigorously  with 
chlorine,  hydrogen  chloride  and  free  bromine  being  produced, 
2HBr  +  C12  -» 2HC1  +  Br2.  What  are  the  relative  volumes  (p. 
150)? 

Chemical  Properties  of  Hydrobromic  Acid  HBr9  Aq.  —  The 

solution  of  the  hydrogen  bromide  in  water  is  an  active  acid  (cf. 
p.  52).  It  conducts  electricity  extremely  well.  In  contact  with 
certain  metals,  and  with  oxides  of  metals  and  hydroxides  of  metals, 
it  behaves  exactly  like  hydrochloric  acid  (p.  146) .  In  the  first  case, 
hydrogen  is  set  free  and  the  bromide  of  the  metal  produced.  In 
the  other  two  cases,  water  and  the  bromides  of  the  metals  are 
produced.  For  example:  Zn(OH)2  +  2HBr  —>  ZnBr2  +  2H2O. 
Oxidizing  agents  set  bromine  free  from  hydrobromic  acid,  even 
sulphuric  acid,  which  does  not  act  upon  hydrochloric  acid,  being 
able  to  do  this  (p.  196).  Chlorine  dissolved  in  water  displaces 
bromine  from  hydrobromic  acid  and  from  soluble  bromides  with 
ease  (test  for  bromides). 

IODINE  I2 

Occurrence.  —  Iodine  occurs  in  sea-water,  about  one-fifth  of 
it  in  algae  and  four-fifths  in  organic  compounds.  Certain  species 
of  sea- weed,  known  in  Scotland  as  kelp  and  in  Normandy  as  varec, 
remove  it  from  the  water.  The  ash  of  the  sea-weed  sometimes 
contains  as  much  as  two  per  cent,  or  even  more.  The  other  chief 
source  of  iodine  is  in  Chile  saltpeter  (mainly  NaNO3),  in  which  it 
is  present  in  the  form  of  about  0.2  per  cent  of  sodium  iodate 
NaI03  and  sodium  iodide.  Most  of  the  iodine  of  commerce  is 
obtained  from  this  source  and  only  a  little  from  sea-weed.  The 
largest  proportion  of  iodine  in  the  human  body  is  in  the  thyroid 


IODINE  199 

gland.  In  diseases  like  goitre  and  cretinism,  where  the  thyroid 
is  ill-developed,  injection  of  a  substance  called  iodothyrine,  ex- 
tracted from  sheep's  thyroids,  produces  marked  improvement. 

Preparation.  —  1.  In  factories  where  the  iodine  is  extracted 
from  sea- weed,  the  latter  is  carbonized  in  retorts  and  sodium  iodide 
is  extracted  with  water  from  the  residue.  This  is  then  treated 
with  manganese  dioxide  and  sulphuric  acid.  The  quantity  of 
manganese  dioxide  is  carefully  measured  so  as  to  be  just  sufficient 
to  set  free  the  iodine  contained  in  the  liquid,  without  proceeding 
farther  to  the  liberation  of  the  chlorine  which  it  contains  in  much 
larger  amounts.  When  the  mixture  is  heated,  the  iodine  passes  off 
in  the  form  of  vapor,  and  is  condensed  in  a  suitable  receiver. 
The  action  (cf.  pp.  157,  194)  is: 

2NaI  +  Mn02  +  3H2S04  ->  MnS04  +  2NaHS04  +  2H2O  +  I2. 

2.  In  France  the  treatment  is  similar,  excepting  that  chlorine  is 
used  to  liberate  the  iodine  in  the  last  stage  (2NaI+Cl2-+2NaCl+I2). 
The  quantity  is  adjusted  so  that  excess  may  not  be  employed. 
The  iodine,  being  insoluble,  forms  a  dense  precipitate  and,  when 
the  liquid  is  pressed  out,  it  remains  behind  in  the  form  of  a 
paste. 

3.  Electricity  could  also  be  used  for  the  decomposition  of  this 
mother-liquor.     The  iodine  is  set  free  at  the  positive  electrode. 

In  all  cases  the  iodine  is  purified  by  distillation  with  a  little 
powdered  potassium  iodide.  It  condenses  in  the  solid  form  di- 
rectly, in  glittering,  black  plates  (sublimed  iodine).  The  distilla- 
tion of  a  solid  body,  when  a  condensation  takes  place  directly  to 
the  solid  form,  is  spoken  of  as  sublimation. 

Physical  Properties.  —  Iodine  (Gk.,  like  a  viokt)  is  a  black, 
solid  substance  (sp.  gr.  5),  exhibiting  large  crystalline  plates  of 
rhombic  form.     It  melts  at  114°,  and  boils  at  184°.    The  vapor  has 
at  first  a  reddish-violet  tint,  and  on  being  more  strongly  heate 
becomes  deep  blue  (see  next  section). 

Iodine  is  very  slightly  soluble  in  water  (about  1  :  6000),  and  the 
solution  has  a  scarcely  perceptible  brown  tint  It  is  much  more 
soluble  in  carbon  disulphide  (p.  12)  and  in  chloroform,  m  which 
it  gives  violet  solutions.  In  alcohol  it  gives  a  solution  which  u 


200  COLLEGE    CHEMISTRY 

brown,  the  iodine  being  in  a  condition  of  feeble  combination,  and 
not  simply  in  solution.  An  aqueous  solution  of  potassium  iodide, 
hydrogen  iodide,  or  any  other  iodide,  has  likewise  the  power  to 
take  up  large  quantities  of  iodine.  Here  the  formation  of  definite 
compounds  (such  as,  KI  +  I2<^KI3),  by  a  reversible  action, 
accounts  for  the  amount  of  iodine  taken  up. 

The  behavior  of  free  iodine  towards  starch  forms  a  distinctive 
test  for  both  substances  (cf.  p.  3).  The  pale-brown  aqueous 
solution,  for  example,  when  added  to  starch  emulsion,  produces  a 
deep-blue  color.  This  blue  substance  is  not  a  chemical  compound. 
The  iodine  is  adsorbed  by  the  starch,  which  is  in  colloidal  suspension 


Chemical  Properties.  —  The  molecular  weight  of  iodine,  ascer- 
tained by  weighing  the  vapor  at  temperatures  from  the  boiling- 
point  up  to  700°,  is  253.8.  The  atomic  weight  being  126.92,  the 
molecule  contains  two  atoms.  Beyond  700°,  the  vapor  diminishes 
in  density  more  rapidly  than  Charles'  law  would  lead  us  to  expect, 
and  at  1700°  the  molecular  weight  has  fallen  to  127  (cf.  p.  117). 
As  the  vapor  is  heated,  a  larger  and  larger  proportion  of  the  mole- 
cules is  broken  up,  until  the  decomposition  has  become  complete. 
As  in  all  cases  of  dissociation,  when  the  vapor  is  cooled  the  atoms 
recombine  to  form  molecules.  This  is  the  most  notable  case  in 
which  we  encounter  both  the  monatomic  and  the  diatomic  forms 
of  the  same  element.  The  heat  given  out  when  the  atoms  reunite 
to  form  the  molecules  is  very  considerable  (21  +±  I2  +  28,500  cal.), 
indicating  that  the  chemical  union  of  two  atoms  of  identical  nature 
may  be  as  vigorous  as  that  of  two  atoms  of  different  chemical 
substances.  The  heat  of  union  of  atomic  hydrogen  (p.  113)  is 
even  greater  (2H  <±  H2  +  90,000  cal.).  In  both  cases,  in  accord- 
ance with  Van't  Hoff's  law  (p.  189),  raising  the  temperature 
increases  the  dissociation,  because  that  is  the  direction  in  which 
heat  is  absorbed. 

Iodine  unites  very  slowly  with  hydrogen,  even  when  heated. 
It  unites  directly  with  some  non-metals  and  with  the  majority  of 
the  metals.  When  phosphorus  is  presented  in  the  yellow  form, 
the  action  takes  place  spontaneously  without  the  assistance  of 
heat.  Both  chlorine  and  bromine  displace  iodine  from  combina- 
tion with  hydrogen  and  the  metals  (2HI  +  Br2  -*  2HBr  +  I,). 


HYDROGEN   IODIDE  201 

The  action  may  be  brought  about  either  with  the  substances  in 
dry  form  or  with  their  aqueous  solutions. 

Iodine  and  its  compounds  are  much  used  in  the  arts  and  medicine. 
Iodine  is  applied,  in  the  form  of  an  alcoholic  solution  (tincture  of 
iodine),  for  the  reduction  of  some  swellings.  It  is  required  in 
making  iodoform  CHI3,  and  the  iodides  of  potassium,  rubidium, 
and  sodium,  which  are  used  in  medicine.  The  emulsion  used  in 
making  photographic  dry-plates  contains  silver  iodide  Agl. 

HYDROGEN  IODIDE  HI 

Preparation.  —  The  direct  union  of  hydrogen  and  iodine  can- 
not be  employed  in  preparing  pure  hydrogen  iodide  (see  below). 

The  action  of  concentrated  sulphuric  acid  upon  potassium  iodide 
is  equally  inapplicable.  In  this  case,  as  in  that  of  hydrogen 
bromide  (p.  196),  the  sulphuric  acid  oxidizes  the  hydrogen  halide 
and  much  free  iodine  and  hydrogen  sulphide  are  formed: 

H2S04  +  SHI  ->  H2S  T  +  4H20  +  4I2  1  . 

The  action  affords  a  rough  test  for  an  iodide  (cf.  pp.  3,  200). 

Powdered  sodium  iodide  and  concentrated  phosphoric  acid  (cf. 
p.  196),  when  warmed,  give  pure  hydrogen  iodide  very  slowly. 

The  best  method  is  one  similar  to  that  described  under  hydrogen 
bromide.     Phosphorus  and  Iodine  unite  directly  to  form  PI3. 
This  is  a  yellow  solid  which  is  violently  hydrolyzed  by  water  and 
gives  phosphorous  acid  and  hydrogen  iodide: 
PI3  +  3H20-»P(OH)3 


If  excess  of  water,  which  dissolves  hydrogen  iodide,  is  avoided,  the 
latter  goes  off  in  a  continuous  stream  in  a  gaseous  condition. 
apparatus  shown  in  Fig.  76  may  be  used.    The  mixture  of  iodine 
and  red  phosphorus  is  placed  in  the  flask  and  the  water  in  the 

funnel.  . 

Still  another  method  of  making  hydrogen  iodide  is  frequen 
employed  when  a  solution  of  the  gas  in  water  is  required,  and 
the  gas  itself.     Powdered  iodine  is  suspended  in  water,  and  hydi 
gen  sulphide  gas  (q.v.)  is  introduced  through  a  tube  in  a  continuous 
stream.     The  iodine  dissolves  slowly  in  the  water,  I,  (solid)  *±I, 
(dslvd),  and  acts  upon  the  hydrogen  sulphide,  which  likewise  du 


202  COLLEGE    CHEMISTRY 

solves,  H2S  (gas)  ^±  H2S  (dslvd).  Sulphur  separates  in  a  fine 
powder,  S  (dslvd)  <=*  S  (solid),  and  hydrogen  iodide  is  formed  in 
accordance  with  the  equation: 

H2S  +  I2  -*  2HI  +  S  | . 

This  action  takes  place,  however,  only  in  presence  of  water,  al- 
though the  water  does  not  appear  in  the  equation.  The  solution 
is  freed  from  the  deposit  of  sulphur  by  nitration,  and  may  be  con- 
centrated to  57  per  cent  of  hydriodic  acid  by  distilling  off  the  water. 

Physical  Properties.  —  Hydrogen  iodide  is  a  colorless  gas  with 
a  sharp  odor.  Its  molecular  weight  is  128,  and  it  is  therefore  much 
heavier  than  air,  the  average  weight  of  whose  molecules  is  28.955 
(p.  101).  It  is  a  nonconductor  of  electricity,  both  in  the  gaseous 
and  in  the  liquefied  conditions.  It  is  exceedingly  soluble  in  water, 
so  that  at  0°  ten  grams  of  water  will  absorb  ninety  grams  of  the  gas, 
giving  a  90  per  cent  solution.  The  behavior  of  this  solution  is  simi- 
lar to  that  of  hydrogen  chloride  and  hydrogen  bromide  (cf.  p.  145). 
The  mixture  of  constant  boiling-point  distils  over  at  127°  (at  760 
mm.),  and  contains  57  per  cent  of  hydrogen  iodide. 

Chemical  Properties.  —  Hydrogen  iodide  is  the  least  stable  of 
the  hydrogen  halides.  When  heated  it  begins  visibly  to  decompose 
into  its  constituents  at  180°.  On  account  of  the  ease  with  which 
it  parts  with  the  hydrogen  which  it  contains,  it  can  be  burned  in 
oxygen  gas,  4HI  +  02  — »  2H20  +  2I2.  When  the  gas  is  mixed 
with  chlorine,  a  violent  chemical  change,  accompanied  by  a  flash  of 
light,  occurs,  the  iodine  is  set  free,  and  hydrogen  chloride  is  pro- 
duced, C12  +  2HI  — >  2HC1  +  I2.  Bromine  vapor  will  similarly 
displace  the  iodine  from  hydrogen  iodide. 

Chemical  Properties  of  Hydriodic  Acid  HI,  Aq.  —  In  most 
respects  the  aqueous  solution  behaves  exactly  like  hydrochloric 
and  hydrobromic  acids.  With  oxidizing  agents,  for  example,  such 
as  manganese  dioxide,  it  gives  free  iodine,  just  as  the  others  (p.  158) 
give  free  chlorine  and  bromine,  respectively.  Here,  however,  the 
oxidation  is  so  much  more  easily  carried  out,  that  it  is  slowly 
effected  by  atmospheric  oxygen,  so  that  hydriodic  acid  left  exposed 
to  the  air  gradually  becomes  brown  (02  +  4HI  — >  2H20  +  2I2). 


HYDROGEN   IODIDE  203 

Although  the  dry  gas  is  not  an  acid,  the  solution  has  all  the  ordi- 
nary properties  of  this  class  of  substances  (cf.  p.  52).  The  hydro- 
gen may  be  displaced  by  metals  like  zinc  and  magnesium  (p.  60). 
The  acid  interacts  with  oxides  and  hydroxides,  forming  iodides  and 
water  (p.  146). 

The  Direct  Union  of  Hydrogen  and  Iodine.  —  The  union  of 
hydrogen  and  iodine,  giving  hydrogen  iodide,  is  a  reversible  re- 
action: 

2HI  <=±  H2  +  I2. 

That  is  to  say,  whether  we  charge  a  tube  with  hydrogen  iodide, 
or  with  an  equal  amount  of  the  elements  in  the  correct  proportions 
by  weight,  if  we  place  both  tubes  in  a  bath,  and  keep  them  thus  at 
the  same  temperature,  the  contents  of  the  tubes  will  after  a  time 
be  identical  (p.  177) .  At  283°,  there  will  be  82  per  cent  of  the  com- 
pound, and  18  per  cent  of  the  uncombined  elements.  At  508°  the 
proportions  will  be  76  per  cent  and  24  per  cent,  respectively. 

The  proportion  of  the  elements  increases  with  rise  in  tempera- 
ture because  the  dissociation  absorbs  heat  (p.  189). 

At  any  one  temperature,  say  283°,  the  equilibrium  point  can 
be  displaced  in  either  direction  (p.  181).  If  we  introduce  some 
additional  hydrogen  (or  iodine),  without  enlarging  the  tube,  thus 
increasing  the  concentration  of  the  hydrogen  (or  iodine),  more 
than  82  per  cent  of  the  compound  is  formed.  If,  instead,  we  let 
one  end  of  the  tube  project,  and  cool  this  end,  the  iodine  con- 
denses to  solid  form,  while  the  other  two  substances  remain 
gaseous.  This  lowers  the  concentration  of  the  iodine  in  the 
gaseous  mixture,  and  lowers  the  speed  and  force  of  the  union  of 
the  elements.  It  does  not  affect  the  tendency  to  dissociation  of  the 
compound  molecules,  but,  since  it  interferes  with  the  formation 
of  more  of  them,  it  enables  the  dissociation  to  proceed  to  practical 
completion.  The  condensation  of  the  iodine  is  essentially  like  a 
precipitation  (pp.  144,  186). 

This  reaction  illustrates  very  clearly  the  way  in  which  the  prog- 
ress of  a  reversible,  chemical  action  is  controlled  by  mechanical 
causes.  It  shows  also  why  we  do  not  prepare  the  compound  by 
uniting  the  elements:  (1)  Since  the  elements  interact  as  gases, 
very  bulky  apparatus  would  be  required  to  prepare  any  consider- 


204 


COLLEGE    CHEMISTRY 


able  quantity;  (2)  the  union  is  very  slow,  taking  many  hours  at 
283°;  (3)  it  is  incomplete,  at  best,  and  we  obtain  a  mixture,  and 
not  a  pure  substance. 

Note  that,  removing  one  product  is,  in  general,  more  effective 
than  increasing  the  concentration  of  one  of  the  interacting  sub- 
stances. The  concentration  of  one  product  can  be  reduced  to 
zero.  To  achieve  the  same  effect  by  adding  an  interacting  sub- 
stance, the  concentration  of  the  latter  would  have  to  be  raised  to 
infinity,  which  is  impossible. 

FLUORINE  F2. 

The  discussion  of  this  element  should  logically  have  preceded 
that  of  chlorine,  since  it  is,  of  all  the  members  of  the  halogen  family, 
the  most  active.  Chlorine  was  taken  up  first,  however,  because 
its  compounds  are  more  familiar.  Fluorine  is  found  in  nature 
chiefly  in  the  mineral  fluorite,  calcium 
fluoride  CaF2  and  in  cryolite,  a  double 
fluoride  of  aluminium  and  sodium  3NaF, 
A1F3. 

Preparation.  —  When  a  solution  of 
hydrofluoric  acid  is  heated  with  man- 
ganese dioxide,  oxidation  does  not  occur 
and  free  fluorine  is  not  produced.  Until 
recently  all  efforts  to  isolate  the  element 
failed.  It  was  perfectly  understood  that 
the  reason  of  these  failures  lay  in  the 
greater  chemical  activity  of  fluorine,  which 
made  it  more  difficult  of  separation  from 
any  state  of  combination  than  the  other 
halogens.  Its  preparation  was  finally 
achieved  by  Moissan  (1886)  by  the  de- 
composition of  anhydrous  hydrogen  fluo- 
ride, which  is  liquid  below  19°,  by  means  of  electricity.  The 
apparatus  (Fig.  77)  is  made  of  copper,  which,  after  receiving  a  thin 
coating  of  the  fluoride,  is  not  further  affected.  To  reduce  the 
tendency  to  chemical  union,  the  whole  is  immersed  in  a  bath  giving 
a  temperature  of  —  23°.  The  electrodes  are  made  of  an  alloy  of 
platinum  and  iridium,  which  is  the  only  material  that  can  resist 


FIG.  77. 


HYDROGEN   FLUORIDE  205 

the  action  of  the  fluorine.  Hydrogen  fluoride,  like  other  hydrogen 
halides,  is  a  nonconductor  of  electricity,  and  a  small  quantity  of 
potassium-hydrogen  fluoride  KHF2  has  to  be  added  to  enable  the 
current  of  electricity  to  pass.  The  fluorine  is  set  free  at  the  posi- 
tive electrode,  and  hydrogen  appears  at  the  negative.  The  U-tube 
is  closed,  after  the  introduction  of  the  hydrogen  fluoride,  by  means 
of  blocks  made  of  calcium  fluoride,  which  is  naturally  unable 
further  to  enter  into  combination  with  fluorine.  For  the  reception 
and  examination  of  the  fluorine  gas,  other  copper  tubes  can  be 
screwed  on  to  the  side  neck  of  the  apparatus,  and,  when  necessary, 
small  windows  of  calcium  fluoride  can  be  provided. 

Physical  Properties.  —  Fluorine  is  a  gas  whose  color  is  like 
that  of  chlorine,  but  somewhat  paler.  Its  density  (38)  shows  that 
the  molecule  is  diatomic  (F2).  The  gas  is  the  most  difficult  of  the 
halogens  to  liquefy.  The  liquid  boils  at  — 186°. 

Chemical  Properties.  —  Fluorine  unites  with  every  element, 
with  the  exception  of  oxygen,  chlorine,  nitrogen,  and  the  members 
of  the  helium  family,  and  in  many  cases  does  so  with  such  vigor 
that  the  union  begins  spontaneously  without  the  assistance  of 
external  heat.  Dry  platinum  and  gold  are  the  elements  least 
affected.  It  explodes  with  hydrogen  at  the  ordinary  temperature, 
without  the  assistance  of  sunlight.  On  the  introduction  of  a  drop 
of  water  into  a  tube  of  fluorine,  the  oxygen  of  the  water  (vapor) 
is  instantly  displaced  by  fluorine,  and  the  vessel  is  filled  with  the 
deep-blue  gas,  ozone:  3F2  +  3H20  ->  3H2F2  +  O3. 

Fluorine  displaces  the  chlorine  in  hydrogen  chloride  as  easily  as 
chlorine  in  turn  displaces  bromine  or  iodine. 

HYDKOGEN  FLUORIDE  H2F2 

Preparation.  —  Pure,  dry  hydrogen  fluoride  is  best  made  by 
heating  potassium-hydrogen  fluoride,  2KHF2  ^  K2F2  +  H2F2  f  . 
For  ordinary  purposes,  however,  the  preparation  of  an  aqueous 
solution  is  the  ultimate  object.  Usually  powdered  calcium  fluoride 
is  treated  with  concentrated  sulphuric  acid,  and  the  mixture  dis- 
tilled in  a  retort  of  platinum  or  lead : 

CaF2  +  H2S04  <^  CaS04  +  H2F2 1 . 


206  COLLEGE    CHEMISTRY 

The  hydrofluoric  acid  passes  over  and  is  caught  in  distilled  water. 
The  aqueous  solution  thus  obtained  has  to  be  kept  in  vessels  made 
of  lead,  rubber,  or  paraffin,  as  glass  interacts  with  the  acid  with 
great  rapidity  (see  below). 

Physical  Properties.  —  Hydrogen  fluoride  is  a  colorless  liquid, 
boiling  at  19.4°.  It  mixes  freely  with  water  and,  on  distillation,  an 
acid  of  constant  boiling-point  (120°  at  760  mm.)  containing  35  per 
cent  of  hydrogen  fluoride  is  obtained.  The  weight  of  22.4  liters 
of  the  vapor  varies  from  20  g.  at  90°  and  above,  to  51  g.  at  26°. 
At  90°,  therefore,  the  formula  is  HF  and  at  26°  probably  a  mixture 
of  H2F2(40)  and  H3F3(60).  Since  HF  is  the  only  form  which  per- 
sists through  a  range  of  temperature,  we  say  this  substance  shows 
association  at  lower  temperatures.  Water  is  spoken  of  as  an 
associated  liquid  —  the  vapor  being  pure  H2O,  but  the  liquid  a 
mixture  of  this  along  with  (H20)2  and  (H20)a  (p.  138). 

Chemical  Properties   of  Hydrofluoric  Acid  HQF29  Aq.  — 

Metals  like  zinc  and  magnesium  interact  with  hydrofluoric  acid 
with  evolution  of  hydrogen  (p.  60).  The  action  is  less  violent  than 
with  other  halogen  acids.  The  acid  interacts  with  oxides  and 
hydroxides,  forming  fluorides  (p.  146).  The  chief  difference  in  this 
respect  which  it  exhibits,  when  compared  with  the  other  halogen 
acids,  is  one  which  leads  us  to  assign  to  it  the  formula,  H2F2.  We 
may  displace  either  one  or  both  the  hydrogen  atoms  in  the  molecule 
with  a  metal.  Thus,  one  of  the  commonest  salts  of  hydrofluoric 
acid  is  potassium-hydrogen  fluoride,  or  the  acid  fluoride  of  potas- 
sium KHF2,  mentioned  above.  In  this  respect  the  acid  resembles 
sulphuric  acid  and  other  acids  containing  more  than  one  replace- 
able hydrogen  unit. 

The  most  remarkable  property  of  hydrofluoric  acid  depends  on 
the  great  tendency  which  fluorine  has  to  unite  with  silicon,  forming 
the  gaseous  silicon  tetrafluoride.  Glass  (q.v.)  is  essentially  a  mix- 
ture of  silicates  of  calcium  and  sodium,  with  excess  of  silica  (sand) 
Si02,  and  is  rapidly  decomposed  by  hydrofluoric  acid: 

CaSi03  +  3H2F2  -»  SiF4  T  +  CaF2  +  3H20, 
Si02  +  4H2F2  ->  SiF4  f  +  2H20. 

In  all  other  silicates,  fluorine  is  substituted  (p.  162)  for  oxygen 


THE   HALOGENS   AS   A   FAMILY 


207 


according  to  the  same  plan.  The  silicon  tetrafluoride  SiF4  is  a  gas. 
The  fluorides  of  calcium  and  sodium  are  solid  and  crumble  away  or 
dissolve.  Thus  the  glass  is  completely  disintegrated.  The  vapor 
of  hydrofluoric  acid,  generated  in  the  way  described  above  from 
calcium  fluoride  in  a  lead  dish,  is  used  for  etching  glass.  The  sur- 
face of  the  glass  is  covered  with  paraffin  to  protect  it  from  the  action 
of  the  vapor,  and  with  a  sharp  instrument  portions  of  this  paraffin 
are  removed  where  the  etching  effect  is  desired..  The  vapor  gives 
a  rough  surface  where  it  encounters  the  glass  (test  for  a  fluoride). 
In  this  way,  the  graduation  on  thermometers,  burettes,  and  other 
pieces  of  apparatus,  is  marked.  The  aqueous  solution  makes 
smooth  depressions  on  the  surface  of  glass.  It  is  used  for  removing 
sand  from  metal  castings  and  for  cleaning  the  exteriors  of  buildings 
of  granite  and  sandstone. 

THE  HALOGENS  AS  A  FAMILY 

The  most  noticeable  fact  is  that,  if  we  arrange  the  halogens  in 
order  in  respect  to  any  one  property,  chemical  or  physical,  the  other 
properties  will  be  found  to  place  them  in  the  same  order.  In  the 

table  the  sixth  column  contains  the  weight  of  the  element  dissolv- 
ing in  100  c.c.  of  water  (15°).  The  last  column,  cal.  KX,  gives i  the 
heat  of  formation  of  one  gram-molecule  of  the  potassium  halide. 


Element. 

Atomic 
Weight. 

State. 

Boiling- 
point. 

Color. 

Solubility. 

Cal.  KX. 

Fluorine    .   . 
Chlorine    .    . 
Bromine    .   . 
Iodine    .   .   . 

19.0 
35.5 
79.9 
126.9 

gas 
gas 
liquid 
solid 

-187° 
-  34° 
+  59° 
184° 

yellow 
yellow 
brown 
violet 

"7*2" 

3.2 
0.015 

118,100 
104,300 
95,100 
80,100 

It  will  be  seen  that,  as  the  atomic  weight  increases,  the foiling 
point  (b.-p.)  rises,  the  color  deepens,  the  solubdity  dinumshes  and 
the  heat  of  union  with  potassium  becomes  smaller     The  vigor 
with  which  the  halogens  unite  with  hydrogen  and  the  metals 
greatest  with  fluorine  and  diminishes  progressively  untd  we  reach 
iodine     We  shall  see  later  that  the  affinity  for  oxygen,  on  the 
other  hand,  increases  as  we  pass  from  fluorine  to  >°dine. 
Although  showing  different  degrees  of  activity,  the  halogens  are 


208  COLLEGE    CHEMISTRY 

closely  alike  in  chemical  nature.  That  is,  the  relations  (p.  163) 
they  show  when  in  combination  are  similar.  When  united  with 
hydrogen  and  the  metals,  they  are  all  univalent.  In  their  oxygen 
compounds,  however,  they  exhibit  a  higher  valence.  Their  oxides 
interact  with  water  to  give  acids,  and  they  are  therefore  non- 
metals  (p.  94).  They  are  strongly  electro-negative  (pp.  55,  194), 
as  non-metals  all  are.  Their  hydrides,  when  dissolved  in  water, 
are  all  active  acids.  This,  and  their  valence,  distinguish  the 
halogen  family  from  other  groups  of  non-metals.  Thus,  oxygen 
and  sulphur  are  bivalent  (and  the  latter  sexivalent  also),  and  the 
hydrides  of  oxygen  (H20  and  H2O2)  and  of  sulphur  (H2S)  are  very 
feeble  acids. 

Order  of  Activity  of  the  Non-Metals.  —  The  way  in  which 
chlorine  displaces  bromine  and  iodine  from  bromides  (p.  194)  and 
iodides  (p.  199),  and  bromine,  in  turn,  displaces  iodine  suggests  an 
order  of  activity  for  non-metals.  It  was  noted  that  oxygen  dis- 
places iodine  from  hydriodic  acid  (p.  202)  and  that  iodine  displaces 
sulphur  from  hydrogen  sulphide  (and  all  other  sulphides).  The 
order  is,  therefore,  F,  Cl,  Br,  0,  I,  S. 


COMPOUNDS  OF  THE  HALOGENS  WITH  EACH  OTHER 

Iodine  unites  directly  with  chlorine  to  form  two  compounds. 
The  more  familiar  one  is  a  red  crystalline  substance,  iodine  mono- 
chloride  IC1.  Another  compound,  IC13,  is  made  by  the  use  of 
excess  of  chlorine.  Iodine  unites  with  bromine  to  form  the  com- 
pound IBr,  while  a  compound  with  fluorine,  IF5  is  supposed  to 
exist.  None  of  these  compounds  are  particularly  stable,  and  some 
of  them  decompose  easily. 

Exercises.  —  1.  What  impurities  is  commercial  iodine  likely  to 
contain?  In  what  way  does  heating  with  potassium  iodide  (p.  199) 
free  it  from  these? 

>    2.   Classify  all  the  chemical  actions  in  this  chapter  according  as 
they  belong  to  one  or  other  of  the  ten  kinds  (p.  166). 

3.  What  are  the  relative  volumes  of  the  gases  in  the  interaction 
of  chlorine  with  hydrogen  bromide  (p.  198),  and  hydrogen  iodide 
(p.  202),  respectively? 


THE   HALOGENS   AS   A   FAMILY  209 

4.  Tabulate,  more  fully  and  specifically  than  is  done  in  the  sec- 
tion on  "The  Halogens  as  a  Family,"  (a)  the  physical  properties, 
(6)  the  chemical  properties,  (c)  the  chemical  relations,  of  the  mem- 
bers of  this  group. 

5.  Construct  the  equation  on  p.  199  by  the  use  of  partial      J 
equations  as  in  the  example  on  p.  195. 

6.  What  are  the  relative  volumes  of  fluorine  and  ozone  in  the 
action  of  the  former  upon  water  (p.  205)? 

7.  What  relative  volumes  of  chlorine  and  iodine  vapor  must  be 
taken  to  make  the  two  chlorides  of  iodine  (p.  208),  respectively? 

8.  At  a  given  temperature,  would  increasing  the  pressure  in  a 
mixture  of  hydrogen  and  bromine  vapor  render  the  union  more  or 
less  complete?    Is  the  action  more  complete  at  a  high  or  at  a  low 
temperature? 


CHAPTER  XVI 
DISSOCIATION  IN   SOLUTION 

THE  employment  of  interacting  substances  in  the  form  of  solu- 
tions is  so  constant  in  chemistry,  and  the  reasons  for  this  are  so 
cogent,  that  we  must  now  resume  the  discussion  of  this  subject 
(c/.  p.  121). 

The  present  chapter  will  be  devoted  to  giving  the  proofs  that 
the  molecules  of  acids,  bases,  and  salts,  in  aqueous  solutions,  are 
actually  dissociated  into  parts  by  the  solvent.  This  will  be  shown 
by  consideration,  successively,  of  certain  peculiarities  in  the 
chemical  behavior,  in  the  freezing-points  and  in  the  boiling-points 
of  the  solutions  of  these  substances.  We  shall  see  that  these  parts 
coincide  in  composition  with  the  radicals. 

Some  Characteristic  Properties  of  Acids,  Bases,  and  Salts, 
Shown  in  Aqueous  Solution.  —  Acids  all  contain  hydrogen 
(p.  53).  In  aqueous  solution,  if  soluble,  they  are  sour  in 
taste,  they  turn  blue  litmus  red,  and  their  hydrogen  is  displaced 
by  certain  metals  (p.  53),  and  has  the  properties  of  a  radical. 
By  the  last  statement  is  meant  that  it  very  readily  exchanges 
places  with  other  radicals  in  reversible  double  decompositions  (p. 
147) .  Amongst  the  acids  mentioned  have  been :  hydrochloric  acid 
HC1,  sulphuric  acid  H2S04,  hypochlorous  acid  HC10,  acetic  acid 
HCC^CHs.  Many  other  bodies,  like  sugar,  kerosene,  and  alcohol, 
contain  hydrogen  also,  but  not  one  of  them  shows  all  of  these 
properties. 

Again,  all  salts  are  made  up  of  two  radicals,  and  the  reversible 
double  decompositions  into  which  they  enter  with  acids,  bases, 
and  other  salts,  consist  in  exchanges  of  these  radicals.  Other 
substances  may  include  the  same  combinations  of  atoms,  but  in 
their  actions  these  groupings  are  often  disregarded.  Thus,  sodium 
chloride  NaCl  and  silver  nitrate  AgNOa  exchange  radicals  com- 
pletely (p.  147)  and,  in  dilute  solution,  hydrogen  chloride  and 

210 


DISSOCIATION   IN    SOLUTION  211 

sodium-hydrogen  sulphate  do  so  partially  (p.  143).  But  sodium 
chloride  and  nitroglycerine  C3H6(N03)3  do  not  interact  at  all.  ^  The 
latter  is  not  a  salt,  although  it  contains  the  same  proportion  of 
nitrogen  to  oxygen  as  does  any  nitrate. 

All  bases  contain  hydroxyl  OH  as  a  radical,  combined  with  some 
positive  radical.  Potassium  hydroxide  KOH  is  soluble  and  active, 
zinc  hydroxide  Zn(OH)2  and  many  others,  however,  are  insoluble. 
Bases  all  exchange  radicals  readily  in  double  decomposition  with 
salts  and  acids.  Other  substances,  like  alcohol  C2H5OH,  may 
contain  hydroxyl,  but  do  not  interact  readily  with  salts  like  NaCl, 
and  are  not  bases. 

The  Influence  of  Water  and  Other  Solvents.  —  It  is  chiefly 
in  aqueous  solution  that  these  special  properties  of  acids,  bases,  and 
salts  become  apparent.  Their  behavior  is  often  quite  different  in 
the  absence  of  this  solvent.  If,  for  example,  we  mix  together  dry 
ammonium  carbonate  (NH4)2C03  and  partially  dehydrated,  solid 
cupric  nitrate  Cu(N03)2,  and  apply  heat,  a  violent  interaction 
begins.  An  immense  cloud  of  smoke  and  gas  is  thrown  out  of  the 
tube,  and  the  substance  remaining  is  either  black,  or  reddish,  in 
parts,  according  to  the  proportions  of  the  substances  employed. 
The  residue  contains  cupric  oxide,  and  sometimes  red  cuprous 
oxide  Cu20.  The  gas  is  tinged  red  by  the  presence  of  nitrogen 
tetroxide  N02,  while  a  more  careful  examination  would  show  that 
it  contained  carbon  dioxide,  nitrogen,  nitrous  oxide  N20,  wate: 
vapor,  and  perhaps  still  other  products. 

The  contrast,  when  these  substances  are  dissolved  in  water  before 
being  brought  in  contact  with  one  another,  is  very  great.  A  pale- 
green  precipitate  is  formed  at  once,  and  rapidly  settles  out  On 
examination,  this  turns  out  to  be  a  carbonate  of  copper  (basic), 
while  evaporation  of  the  solution  furnishes  us  with  ammonium 
nitrate.  There  are  only  two  main  products,  and  the  essenti* 
part  of  the  action  in  solution  may  be  represented  by  the  equat 

(NH4)2C03  +  Cu(N03)2  -» CuC03 1  +  2NH4N03. 
In  the  interaction  between  the  dry  substances  the  molecules  are 
completely  disintegrated,  the  whole  change  is  very  complex,  a 
it  takes  a^ood  deal  of  time.     In  the  action  in  water  -  heating  i 
reauired,  the  substances  are  neatly  broken  apart,  certain  groups 


212  COLLEGE    CHEMISTRY 

of  atoms,  which  we  call  radicals,  are  transferred  as  wholes  from 
one  state  of  combination  to  another,  and  the  rearrangement  takes 
place  instantaneously  in  a  machine-like  manner.  Contrasts  like 
this  between  the  interactions  of  anhydrous  and  dissolved  bodies 
are  very  common. 

Many  compounds,  however,  do  not  show  any  change  in  be- 
havior when  dissolved  in  water.  Sugar,  for  example,  is,  as  a  rule, 
more  readily  acted  upon  in  the  absence  of  any  solvent.  Then 
again,  while  water  is  not  the  only  solvent  which  has  the  effect  we 
have  just  described,  the  majority  of  solvents,  if  they  affect  chemi- 
cal change  at  all,  simply  retard  it.  Thus  the  union  of  iodine  and 
phosphorus  in  the  absence  of  a  solvent  takes  place  spontaneously 
with  a  violent  evolution  of  heat.  When  the  elements  are  dissolved 
in  carbon  bisulphide,  before  being  mixed,  the  action  is  much  milder, 
although  the  product  is  the  same  (phosphorus  tri-iodide).  The 
diminution  in  the  concentration  of  the  ingredients  has  decreased 
the  speed  of  the  action  in  the  normal  way  (p.  181).  That  water 
and  some  other  solvents  have  a  specific  influence  tending  to  in- 
crease the  apparent  activity  of  certain  classes  of  substances,  shows 
that  a  special  explanation  of  the  phenomenon  must  be  found. 

Summing  up  these  points  we  see  that  the  peculiarity  of  acids, 
bases,  and  salts  in  aqwous  solution  is  that  the  action  is  complete 
as  soon  as  the  solutions  have  been  mixed,  and  that  each  compound 
always  splits  in  the  same  way.  Thus,  cupric  nitrate  always  gives 
changes  involving  Cu  and  N03  and  never  interacts  so  as  to  use 
CuN2  and  03,  or  CuO2  and  N02,  as  the  basis  of  exchange.  Simi- 
larly, dilute  acids  always  offer  hydrogen  in  exchange,  and  so  nitric 
acid  behaves  as  if  composed  of  H  and  NOs,  and  sulphuric  acid  as 
if  composed  of  2H  and  SO4,  and  never  as  if  made  up  of  HSO  and 
HO3,  or  H2S  and  04.  The  sour  taste  and  the  effect  upon  litmus 
seem  to  be  properties  of  this  easily  separable  hydrogen,  for  they 
are  shown  only  by  acids.  The  result  is  that  we  can  make  a  list  of 
the  units  of  exchange,  such  as  H,  OH,  N03,  CO3,  S04,  Cu,  K,  and 
Cl,  employed  by  acids,  bases,  and  salts  in  their  interactions.  The 
molecule  of  each  compound  of  these  classes  contains  at  least  two 
of  them.  Even  when  these  units  contain  more  than  one  atom, 
their  coherence  is  as  noticeable  within  this  class  of  actions,  as  is 
the  permanence  of  the  atomic  masses  themselves  in  all  actions. 

The  question  raised  in  our  minds  is  whether  solution  in  water 


DISSOCIATION   IN   SOLUTION  213 

alters  the  character  of  the  molecule,  simply  by  producing  a  sort  of 
plane  of  cleavage  in  it  which  creates  a  predisposition  to  a  uniform 
kind  of  chemical  change,  or  whether  it  actually  divides  the  molecules 
into  separate  parts  consisting  of  the  above  units  of  exchange,  and 
leaves  subsequent  chemical  actions  to  occur  by  cross-combination 
of  these  fragments.  The  fact  that  the  dissolved  substances  can  be 
recovered  by  evaporation  of  the  liquid  does  not  demonstrate  that 
they  have  not  been  decomposed  temporarily  while  in  solution. 
The  alteration  which  the  water  produces,  whatever  it  be,  will 
naturally  be  reversed  when  the  water  is  removed.  Since  our 
question  involves  nothing  but  the  counting  of  particles,  the  num- 
ber of  which  would  be  much  greater  in  the  event  that  actual  sub- 
division of  molecules  is  the  explanation,  it  can  be  answered  by  a 
study  of  the  physical  properties  of  solutions.  Several  physical 
properties  can  be  used,  and  they  give  concordant  answers  to  the 
question.  We  shall  confine  ourselves  here,  however,  mainly  to  the 
evidence  furnished  by  the  freezing-points  and  boiling-points  of 
solutions. 

Laws  of  Freezing-Point  Depression.  —  Every  pure  liquid 
has  a  definite  temperature  at  which  it  freezes.  Thus,  pure  water 
freezes  at  0°  and  benzene  at  5.48°.  As  we  have  seen  (p.  134),  how- 
ever, the  presence  of  a  foreign,  dissolved  body  lowers  the  freezing- 
point,  although  the  "ice"  which  separates  usually  consists  of 
crystals  of  the  pure  solvent  only. 

The  depression  in  the  freezing-point  is  directly  proportional  to 
the  weight  of  dissolved  substance  in  a  given  amount  of  the  solvent. 
The  depression  is  inversely  proportional  to  the  amount  of  solvent. 
Thus,  if  we  double  the  concentration  of  the  solution,  the  depression 
in  the  freezing-point  is  doubled.  Thus,  in  one  set  of  experiments, 
solutions  of  sugar  containing  11.4  g.,  22.8  g.,  and  34.2  g.  of  sugar 
to  100  g.  of  water  were  found  to  freeze  at  -0.62°,  -1.24°,  and 
—  1.86°,  respectively. 

Further,  equal  numbers  of  molecules  of  different  solutes  in  the 
same  quantity  of  solvent  give  equal  depressions.  Or,  in  other  words, 
the  depression  is  proportional  to  the  concentration  of  the  molecules 
of  the  solute .  Thus,  solutions  containing  342  g.  of  sugar  C^H^O 
or  46  g.  of  alcohol  C2H60,  or  74  g.  of  methyl  acetate 
*  12  X  12  +  22  X  1  +  11  X  16  =  342. 


214  COLLEGE    CHEMISTRY 

in  1000  g,  of  water,  being  weights  which  contain  equal  numbers  of 
molecules,  show  a  depression  below  the  freezing-point  of  water  of 
about  1.86°  in  each  case.  That  is,  such  solutions  all  freeze  close 
to  —1.86°.  This  depression,  produced  by  a  mole  of  the  solute  hi 
1 1.  of  solvent,  is  called  the  molecular  depression  constant,  and  has 
a  different  value  for  each  solvent.  For  solutions  of  the  same 
molecular  concentration  in  benzene  the  depression  is  4.9°,  in 
phenol  (carbolic  acid)  7.3°.  Combining  these  facts  in  one  ex- 
pression: 

The  observed  depression  1  =         0          Wt.  of  Solute  1000 

in  an  aqueous  solution  J  Mol.  Wt.  of  Solute      Wt.  of  Solvent' 

For  other  solvents,  the  corresponding  value  of  the  depression  con- 
stant must  be  substituted  for  1.86°. 

These  laws  describe  the  facts  most  exactly  when  the  solutions 
are  dilute.  They  hold  only  when  there  is  no  chemical  interaction 
between  solute  and  solvent.  Even  so,  however,  adds,  bases,  and 
salts  dissolved  in  water  present  many  apparent  exceptions  and  must 
be  discussed  separately  (see  below). 

It  will  be  noted  that,  when  the  other  factors  in  the  foregoing 
equation  are  known  or  observed,  the  molecular  weight  of  the  solute 
may  be  determined.  The  fact  makes  possible  the  determination  of 
this  constant  for  substances  which  are  not  volatile  (see  Hydrogen 
peroxide) . 

Abnormal  freezing-Point  Depression:  Dissociation  in 
Solution.  —  The  substances  which  present  the  most  conspicuous 
exceptions  to  the  above  rules  are  acids,  bases,  and  salts  in  aqueous 
solution.  With  most  of  these,  the  depression  produced  is  abnormal ; 
it  is  greater  than  we  should  expect  from  the  concentration  of  the 
solution.  Thus,  in  an  actual  experiment,  two  equi-molar  solu- 
tions were  compared.  One  contained  one  mole  (74  g.)  of  methyl 
acetate,  and  the  other  one  mole  (58.5  g.)  of  sodium  chloride,  each 
dissolved  in  2000  g.  (2  liters)  of  water.  The  freezing-points 
observed  were: 

Pure  water 0.000°        Pure  water 0.000° 

Sol.  of  methyl  acetate  .    -0.970°         Solution  of  salt  ....    -1.678° 

Depression 0.970°         Depression 1.678° 

0.970° 
Excess  depression  by  salt    0 . 708° 


DISSOCIATION   IN   SOLUTION  215 

The  solution  of  methyl  acetate,  as  it  contained  only  0.5  moles  of 
the  solute  per  liter  of  water,  showed,  as  it  should  do,  about  half  the 
average  molecular  depression  (1.86°,  p.  214).  This  is  typical  of 
the  class  of  substances  showing  normal  behavior.  Sugar,  alcohol, 
and  hundreds  of  other  substances,  in  solutions  of  the  same  molar 
concentration,  would  have  given  the  same  value. 

The  freezing-point  of  the  salt  solution,  however,  was  much  lower. 
If  this  solution  had  really  contained  the  same  concentration  of  dis- 
solved molecules  as  the  other  solution,  its  depression  would  like- 
wise have  been  0.970°.  The  number  of  molecules  in  the  solution 
must  therefore  have  been  greater  than  we  should  have  expected 
from  the  number  of  molecules  taken.  In  other  words,  a  portion 
of  the  molecules  of  the  salt  must  have  been  broken  up,  and  the 
excess  depression,  0.708°,  must  have  been  due  to  the  extra  mole- 
cules produced  by  dissociation.  Now  sodium  chloride  molecules 
cannot  give  more  than  two  particles  each,  and  the  depression  is 
proportional  to  the  number  of  particles.  It  follows,  therefore, 
that  J$f,  or  0.732  (73.2  per  cent)  of  the  molecules  were  dissociated: 

(27  per  cent)  NaCl<=±  (Na)  +  (Cl)  (73  per  cent). 

This  result  is  typical  also.  Acids,  bases,  and  salts,  of  which  one 
mole  is  dissolved  in  two  liters  of  water,  are  found  to  give  irregular 
values,  all  more  or  less  in  excess  of  0.970°.  Those  which  contain 
but  two  radicals,  like  sodium  chloride  NaCl  and  potassium  nitrate 
KNO3,  give  values  between  0.970°  and  2  X  0.970°.  Substances 
like  calcium  chloride  Ca(Cl)2  and  sodium  sulphate  (Na)2S04  give 
depressions  approaching  three  times  the  normal  value:  their 
molecules  contain  three  radicals.  The  excess  depression  depends, 
therefore,  upon  the  number  of  particles  which  each  molecule  can 
furnish,  and  upon  the  proportion  of  all  the  molecules  which  is 
dissociated  into  these  fragments. 

In  the  case  of  an  acid,  base,  or  salt,  the  depression  is  not  strictly 
proportional  to  the  concentration.  Thus,  one  mole  of  salt  in  four 
liters  of  water  does  not  give  half  the  depression  of  the  two-liter 
solution  (1.678°  •*•  2  =  0.839°)  but  somewhat  more  (about  0.844°). 
The  same  method  of  calculation  indicates,  therefore,  a  greater 
degree  of  dissociation  (about  79  per  cent)  in  the  more  dilute  solu- 
tion (see  Ionic  equilibrium). 

Acids,  bases,  and  salts,  so  far  as  they  are  soluble  in  materials  like 


216  COLLEGE    CHEMISTRY 

toluene,  benzene,  chloroform,  and  carbon  bisulphide,  exhibit 
simply  normal  depressions  in  these  solvents.  It  appears,  there- 
fore, that,  in  many  solvents,  dissociation  does  not  take  place.  In 
common  experience  it  is  encountered  only  in  solutions  in  water, 
and  in  alcohol. 

Abnormal  Boiling-Point  Elevation.  —  We  have  seen  (p.  135) 
that  342  g.  of  sugar,  or  an  equal  number  of  molecules  of  glycerine 
CaHsOs  (92  g.),  dissolved  in  1000  c.c.  of  water,  will  elevate  the 
boiling  point  from  100°  to  100.52°.  One  molecular  weight  of 
sodium  chloride  (58.5  g.),  however,  will  elevate  the  boiling-point 
of  the  water  0.87°  instead  of  0.52°.  The  effect  is  0.35°,  or  67  per 
cent  greater,  indicating  dissociation  of  this  proportion  of  the  NaCl 
molecules.  In  more  dilute  solutions,  the  elevation  is  relatively 
greater.  Salts  containing  more  than  two  radicals,  like  Ca(Cl)2, 
give  elevations  of  more  than  twice  the  normal  value.  In  solvents 
like  benzene  and  carbon  disulphide,  however,  no  abnormal  eleva- 
tion is  observed  with  any  solute.  The  phenomena  are,  in  fact, 
parallel  with  those  connected  with  the  freezing-point. 

Other  Evidence  of  Dissociation.  —  The  freezing-point  and 
boiling-point  are  only  two  of  four  properties  of  solutions  which  can 
be  used  for  determining  the  numbers  of  molecules  present.  Nu- 
merous measurements  show  that  aqueous  solutions  of  acids,  bases, 
and  salts  have  also  abnormal  osmotic  pressures  (c/.  p.  135).  The 
electrical  conductivity  is  the  fourth  property  which  gives  the 
required  information  (see  Chap.  XVIII).  Now,  when  we  observe 
the  behavior  of  the  same  solution  in  each  of  these  four  ways,  and 
calculate  the  degree  of  dissociation  from  the  result  of  each  measure- 
ment, we  find  that  the  values  obtained  are  usually  identical,  within 
the  limits  of  error  to  which  the  methods  are  liable.  Thus  the  in- 
dications of  dissociation  found  in  the  chemical  behavior  of  acids, 
bases,  and  salts  (pp.  211-213)  are  fully  confirmed  by  a  study  of  the 
physical  properties  of  their  solutions. 

Applications:  The  Constitution  of  Solutions  of  Acids, 
Bases,  and  Salts.  —  The  composition  of  solutions  which  are  nor- 
mal or  abnormal,  in  respect  to  osmotic  pressure,  freezing-point,  and 
boiling-point,  may  be  shown  thus: 


DISSOCIATION    IN    SOLUTION  217 


Solutes. 

Dissolved  in 
Water,  Alcohol, 
etc. 

Dissolved  in 
Toluene,  Chlo- 
roform, etc. 

Acids,  bases,  salts  

Abnormal 

Normal 

Other  substances  

Normal 

Normal 

It  appears  that  water  and  some  other  solvents  have  the  power  of 
decomposing  acids,  bases,  and  salts.  Such  solvents  have,  in  fact, 
an  effect  on  these  materials  that  resembles,  outwardly  at  least,  the 
effect  which  heat  has  on  many  substances  (e.g.,  p.  117),  they  cause 
dissociation:  CaCl2^.(Ca)  +  2(C1). 

In  consequence  of  this,  our  view  of  the  nature  of  an  aqueous  solu- 
tion of  hydrogen  chloride  HC1,  or  common  salt  NaCl,  or  sodium 
hydroxide  NaOH,  or  any  of  the  substances  of  the  classes  which 
these  represent,  may  now  be  stated  in  definite  terms.  Such  a  solu- 
tion contains,  besides  undivided  molecules  of  the  solute,  at  least 
two  other  kinds  of  material,  H,  Na,*  Cl,  OH,  etc.,  which  result  from 
the  breaking  up  of  the  molecules.  We  shall  see  that  these  sub- 
divisions of  the  original  molecules  have  distinct  physical  and  chemi- 
cal properties  of  their  own.  The  descriptions  of  the  "properties" 
of  the  solutions,  as  they  used  to  be  given  in  chemistry,  were  really 
p,  confused  statement  of  the  properties  of  the  different  components 
of  a  mixture. 

The  free  radicals,  of  whose  existence  we  have  thus  become  con- 
vinced, constitute  a  new  set  of  materials  (with  appropriate  names. 
See  p.  236).  Thus  the  hydrogen  radical  of  acids,  although  a  form 
of  uncombined  hydrogen,  differs  totally  from  the  gas  which  is  com- 
posed of  the  same  material.  The  gas  has  no  sour  taste  or  effect 
upon  litmus;  these  are  properties  of  the  free  radical.  The  gas  is 
very  slightly  soluble  in  water,  while  the  hydrogen  radical  exists  as  a 
separate  substance  only  in  solution.  Again,  substances  with  the 
composition  of  the  radicals  N03  and  S04  are  not  known  at  all 
except  in  solutions. 

Exercises.  —  1.  What  depression  in  the  f.-p.  of  water  will  be 
produced  by  dissolving  10  g.  of  bromine  in  1  kg.  of  this  solvent? 

*  The  objection  that  separate  atoms  of  sodium  could  not  remain  free  in 
water,  will  be  disposed  of  later. 


218 


COLLEGE   CHEMISTRY 


2.  What  depressions  in  the  f  .-p.  of  benzene  and  of  phenol  would 
be  produced  by  10  g.  of  bromine  to  1  kg.  of  the  solvent,  if  no 
chemical  action  took  place? 

3.  What  is  the  molecular  depression-constant  of  a  solvent  in 
which  5  g.  of  iodine  in  500  g.  of  the  solvent  lowers  the  f.-p.  0.7°? 

4.  What  is  the  degree  of  dissociation  of  zinc  sulphate,  if  5  g.  of  it 
dissolved  in  125  g.  of  water  produce  a  lowering  of  0.603°  in  the  f.-p.? 

5.  In  a  decinormal  solution,  potassium  chloride  is  86  per  cent 
ionized.     What  is  the  freezing  point  of  this  solution? 


CHAPTER  XVII 
OZONE  AND  HYDROGEN  PEROXIDE 

A^  FRESH,  penetrating  odor,  resembling  that  of  very  dilute 
chlorine,  was  noticed  by  van  Marum  (1785)  near  an  electrical 
machine  in  operation.  Schonbein  (1840)  showed  that  the  odor 
was  that  of  a  distinct  substance,  which  he  named  ozone  (Gk.,  to 
smell),  and  he  discovered  a  number  of  ways  of  obtaining  it.  It  is 
very  questionable  whether  there  is  any  ozone  in  the  air,  excepting 
temporarily  in  the  immediate  neighborhood  of  a  natural  or  artificial 
discharge  of  electricity. 

Preparation  of  Ozone  O3.  —  The  most  satisfactory  way  of 
preparing  ozone  is  to  allow  electric  waves  to  pass  through  oxygen. 
The  apparatus  (Fig.  78)  consists  of  two  co-axial  glass  tubes,  be- 
tween which  the  oxygen  flows.  The  waves  are  generated  by  con- 


FIG.  78. 


necting  an  outer  layer  of  tinfoil  on  the  outer  tube,  and  an  inner 
of  tinfoil  in  the  inner  tube  with  the  poles  of  an  induction  cou. 
^ry,  cold  oxygen,  about  7.5  per  cent  of  the  gas  is  turned  into 

Tone  is  found  in  the  oxygen  generated  bj ^electrolysis ;  of dilute 
.nlnnuric  acid  (p.  55).     It  arises  during  the  slow  oxidation  of 


220  COLLEGE   CHEMISTRY 

when  a  jet  of  burning  hydrogen,  or  an  electrically  heated  loop  of 
platinum  wire  is  immersed  in  liquid  oxygen.  This  method  shows 
that  ozone  is  formed  at  high  temperatures,  and  survives  when 
cooled  suddenly  by  the  liquid  oxygen. 

Physical  Properties  of  Ozone.  —  Ozone  is  a  gas  of  blue  color. 
It  boils  at  —119°,  so  that  when  a  mixture  of  oxygen  and  ozone  is 
led  through  a  U-tube  immersed  in  liquid  oxygen  (—182.5°),  the 
ozone  collects  in  the  tube  as  a  deep-blue  fluid.  Ozone  is  much 
more  soluble  in  water  than  is  oxygen.  At  12°,  100  volumes  of 
water  would  dissolve  50  volumes  of  the  gas  at  one  atmosphere 
pressure. 

Chemical  Properties  of  Ozone.  —  The  density  of  ozone  is 
one-half  greater  than  that  of  oxygen.  Its  molecular  weight  is 
therefore  48,  and  its  formula  O3.  Being  formed  with  absorption 
of  energy,  ozone  is  most  stable  at  very  high  temperatures  (Van't 
Hoff's  law,  p.  188). 

302  +  61,400  cal.  <=±  203. 

When  produced  in  cold  oxygen,  by  energy  from  electric  waves,  it 
decomposes  slowly.  But  this  change,  like  all  others,  is  hastened 
by  raising  the  temperature.  Equilibrium,  with  almost  no  ozone, 
is  reached  instantly  at  250-300°.  Liquid  ozone  sometimes  de- 
composes explosively.  As  the  equation  shows,  three  volumes  of 
oxygen  give  two  of  ozone. 

Ozone  is  a  much  more  active  oxidizing  agent  than  oxygen.  Mer- 
cury and  silver,  which  are  not  affected  by  the  latter,  are  converted 
into  oxides  by  the  former.  Silver  gives  the  peroxide,  Ag2O2,  thus : 

2Ag  +  203  -»  Ag-A  +  202. 

Paper  dipped  in  starch  emulsion  containing  a  little  potassium 
iodide  is  used  as  a  test  for  ozone: 

03  +  2KI  +  H20  -*  02  +  2KOH  +  I2. 

The  iodine  gives  a  deep-blue  color  to  the  starch  (cf.  p.  200).  This 
test,  however,  will  not  distinguish  ozone  from  chlorine  or  hydrogen 
peroxide,  and  may,  therefore,  be  used  only  in  the  absence  of  these 
substances. 


OZONE  AND  HYDROGEN  PEROXIDE          221 

Ozone  also  removes  the  color  from  many  of  the  vegetable  color- 
ing matters  and  artificial  dyes.  It  should  be  understood  that  the 
great  majority  of  the  complex  compounds  of  carbon  are  colorless. 
Even  a  slight  chemical  change,  affecting  only  one  or  two  of  the 
atoms  in  a  complex  molecule,  is  thus  almost  sure  to  give  a  color- 
less or  much  less  strongly  colored  material.  Indigo,  Ci6Hi0N202, 
which  has  a  deep-blue  color,  is  an  example  of  a  vegetable  dye  that 
is  also  made  artificially.  When  ozonized  air  is  bubbled  through  a 
dilute  solution  of  this  dye  (as  indigo-carmine),  the  indigo  is  oxidized 
to  isatin  C8H5N02,  and  the  color  disappears  (see  below). 

Ozone  is  used  commercially  in  bleaching  oils,  waxes,  ivory, 
flour,  and  starch.  It  is  employed  also  for  sterilizing  drinking 
water  in  Petrograd,  Lille,  and  other  cities.  For  this  purpose,  how- 
ever, bleaching  powder  is  less  expensive. 

Oxidizing  Agents,  and  Explanation  of  their  Activity.  — 

When  ozone  turns  into  oxygen  much  heat  is  liberated  (equation, 
above).  Ozone  possesses,  therefore,  much  more  internal  energy 
than  does  oxygen.  On  this  account  it  brings  to  the  task  of  oxidiz- 
ing any  substance  more  energy  than  does  oxygen  itself,  and  is  there- 
fore more  efficient.  Thus,  free  oxygen  does  not  niter  act  in  the 
cold  with  indigo,  or  with  silver  or  potassium  iodide  (see  above), 
while  ozone  oxidizes  them  rapidly. 

The  heats  of  reaction  show  the  difference  very  clearly.  In 
equation  (2),  1800  cal.  is  the  amount  of  heat  which  would  be 
liberated  if  indigo  could  be  oxidized  to  isatin  by  oxygen  gas. 
When  ozone  is  used,  we  obtain,  in  addition,  the  heat  of  decompo- 
sition of  this  substance  (equation  1),  so  that  the  total  heat  liber- 
ated (equation  3),  63,200  cal.,  is  35  times  as  great  as  in  equation 
(2)  where  free  oxygen  is  the  oxidizing  agent: 

203  =  202  (+  20)       +       61,400  cal.         (1) 

CieHioNaOs  +  (20)  =  2C8H5NQ2         +         1800  cal.         (2) 

*     C16H10N202  +  203  =  2C8H5N02  +  2O2  +  63,200  cal.         (3) 

By  similar  reasoning  we  explain  the  superiority  of  potassium  per- 
manganate over  free  oxygen  for  oxidizing  hydrochloric  acid  (p. 
157).  Substances  which  are  more  active  oxidizers  than  is  free 
oxygen  may  be  called  active  oxidizing  agents. 


222  COLLEGE    CHEMISTRY 

It  should  be  noted  that  when  ozone  acts  as  an  oxidizing  agent, 
usually  only  one  of  the  atoms  of  oxygen  in  each  molecule  plays 
this  part,  and  oxygen  gas  is  formed.  This  is  illustrated  in  all  the 
three  examples  cited  in  the  preceding  section. 

Allotropic  Modifications.  —  We  have  seen  that  a  substance 
may  exist  in  more  than  the  three  regular  states,  solid,  liquid,  and 
gaseous.  When  a  simple  substance  shows  more  than  one  form, 
in  the  same  state,  like  oxygen  and  ozone,  we  call  them  allotropic 
modifications. 

HYDROGEN  PEROXIDE  H202 

Hydrogen  peroxide  is  found  in  minute  amounts  in  rain  and  snow. 
It  is  formed  in  small  quantities,  in  a  way  not  at  present  fully  under- 
stood, when  moist  metals,  like  zinc,  lead,  and  copper,  rust. 

Preparation  of  Hydrogen  Peroxide.  —  When  sodium  peroxide 
is  added,  a  little  at  a  time,  to  a  cold  dilute  acid,  hydrogen  peroxide 
is  set  free  and  remains  dissolved  in  the  liquid. 

Na202  +  2HC1  fc;  2NaCl  +  H202. 

When  hydrated  barium  peroxide  (Ba02,8H20)  is  shaken  with 
cold,  dilute  sulphuric  acid  a  similar  action  takes  place: 

Ba02  +  H2SO4  fc?  BaSO4 J  +  H2O2. 

Phosphoric  acid  is  largely  employed  instead  of  sulphuric  acid  in  the 
commercial  manufacture  of  hydrogen  peroxide,  and  great  care  is 
taken  to  precipitate  the  other  products  and  all  impurities  from  the 
solution. 

An  aqueous  solution  is  also  obtained  by  passing  carbon  dioxide 
through  barium  peroxide  suspended  in  water : 

Ba02  +  C02  +  H20  fc?  BaC03  j  +  H2O2. 

Pure  hydrogen  peroxide  is  isolated  from  any  of  these  solutions  by 
distillation  under  reduced  pressure.  To  secure  the  low  pressure, 
the  ordinary  distilling  apparatus  (Fig.  51,  p.  93)  is  made  com- 
pletely air-tight,  and  is  connected  by  a  branch  tube  with  a  water- 
pump.  Hydrogen  peroxide  is  much  less  volatile  than  water,  but 
decomposes  into  water  and  oxygen  violently  at  100°.  Hence  the 


HYDROGEN   PEROXIDE  223 

lower  pressure  is  required  to  make  possible  its  volatilization  at  a 
temperature  below  this  point.  At  26  mm.  pressure,  the  water 
begins  to  pass  off  first  (at  about  27°).  The  last  portion  of  the 
liquid  boils  at  69°  and  is  hydrogen  peroxide. 

By  evaporating  the  commercial  (3  per  cent)  solution  at  70°,  a 
liquid  containing  45  per  cent  of  hydrogen  peroxide  may  be  made 
without  much  loss  of  the  material  by  volatilization. 

Physical  Properties.  —  Hydrogen  peroxide  (100%)  is  a  syrupy 
liquid  of  sp.  gr.  1.5.  It  blisters  the  skin  and,  when  diluted,  has 
a  disagreeable  metallic  taste.  It  has  been  frozen  (m.-p.  —2°). 

Chemical  Properties.  —  Hydrogen  peroxide  (100  per  cent)  is 
very  unstable,  and  decomposes  slowly  even  at  —20°.  The  dilute 
aqueous  solution,  when  free  from  impurities,  keeps  fairly  well. 
The  presence  of  a  trace  of  free  acid  increases  its  stability.  Free 
alkalies  and  most  salts  assist  the  decomposition;  hence  the  neces- 
sity for  purifying  the  commercial  solution.  Addition  of  powdered 
metals,  of  manganese  dioxide,  or  of  charcoal  (contact  action) 
causes  effervescence  even  in  dilute  solutions,  and  oxygen  escapes: 

2H202  4  2H20  +  02. 

Since  the  substance  cannot  be  vaporized,  even  at  low  pressure, 
without  some  decomposition,  its  molar  weight  has  been  determined 
by  the  freezing-point  method.  The  freezing-point  of  a  3.3  per  cent 
solution  in  water  was  -2.03°.  Substitution  of  these  data  in  the 
formula  (p.  214)  gives  31.8  g.  as  the  molar  weight.  Now  the  for- 
mula HO  corresponds  to  a  molar  weight  of  17  and  H202  to  one  of 
34.  It  is  evident,  therefore,  that  the  latter  is  the  correct  formula. 

Hydrogen  peroxide,  in  solution  in  water,  is  a  feeble  acid.  As  an 
acid  it  enters  into  double  decomposition  readily,  and  the  peroxides 
are  salts  with  the  negative  radical  02n  (peroxidates) .  Thus, 
when  hydrogen  peroxide  is  added  to  solutions  of  barium  and 
strontium  hydroxides,  the  hydrated  peroxides  appear  as  crystalline 
precipitates : 

Sr(OH)2  +  H202  *±  2H20  +  Sr02. 

The  precipitation  involves  another  equilibrium:  Sr02  +  8H20  +± 
Sr02,8H20  (solid). 


224  COLLEGE    CHEMISTRY 

The  formation  of  a  beautiful  blue  substance  by  the  action  of 
hydrogen  peroxide  upon  dichromic  acid  is  used  as  a  test.  The 
test  is  carried  out  by  adding  a  drop  of  potassium  dichromate  to  an 
acidulated  solution  of  the  peroxide.  The  acid  interacts  with  the 
dichromate,  giving  free  dichromic  acid: 

H2S04  +  K2Cr207  <=*  H2Cr207  +  K2S04. 

The  blue  substance,  which  is  very  unstable  and  quickly  decom- 
poses, is  a  perchromic  acid.  A  blue,  crystalline  perchromic  acid 
(HO)4Cr(OOH)3,  which  decomposes  above  —30°,  has  been  pre- 
pared. The  blue  substance  has  the  property,  unusual  in  inor- 
ganic compounds,  of  dissolving  much  more  readily  in  ether  than 
in  water.  It  is  also  much  less  unstable  when  removed  from  the 
foreign  materials  in  the  aqueous  solution.  Hence  the  test  is 
rendered  more  delicate  by  extracting  the  solution  with  a  small 
amount  of  ether.  In  the  ethereal  layer  the  color  of  the  com- 
pound is  more  permanent,  as  well  as  more  distinctly  visible  on 
account  of  the  greater  concentration. 

Hydrogen  peroxide  is  a  much  more  active  oxidizing  agent  than  is 
free  oxygen.  This  would  be  expected  from  the  fact,  that  it  con- 
tains so  much  more  internal  energy  than  the  water  and  oxygen 
into  which  it  decomposes  (p.  223),  that  23,100  cal.  are  liberated  in 
the  decomposition  of  one  mole.  Thus,  it  liberates  iodine  from 
hydrogen  iodide,  an  action  which,  in  presence  of  starch  emulsion 
(cf.  p.  200),  is  used  as  a  test  for  its  presence: 

2HI  +  H202  -»  2H20  +  I2. 

It  converts  sulphides  into  sulphates.  The  white  lead  (q.v.)  used  in 
paintings  is  changed  by  the  hydrogen  sulphide  in  the  air  of  cities  to 
black  lead  sulphide:  Pb3(OH)2(C03)2  +  3H2S  -»  3PbS  +  4H20  + 
2CO2.  This  may  be  oxidized  to  white  lead  sulphate  by  means  of 
hydrogen  peroxide: 

PbS  +  4H202  ->  PbS04  +  4H20, 

and  in  this  way  the  original  tints  of  the  picture  may  be  practically 
restored.  Organic  coloring  matters  are  changed  into  colorless  sub- 
stances by  an  action  similar  to  that  of  ozone  (cf.  p.  221).  Hence 
hydrogen  peroxide  is  used  for  bleaching  silk,  feathers,  hair,  and 
ivory,  which  would  be  destroyed  by  a  more  violent  agent.  The 


HYDROGEN   PEROXIDE  225 

products  of  its  decomposition,  being  water  and  oxygen  only,  are 
harmless,  and,  on  this  account,  it  is  used  in  disinfecting  (destroy- 
ing organisms  in)  sores,  and  as  a  throat  wash. 

Hydrogen  peroxide  exercises  the  functions  of  a  reducing  agent  in 
special  cases,  also.     Thus,  silver  oxide  is  reduced  by  it  to  silver: 

H202  -»  2Ag  +  H20  +  02. 


A  solution  of  potassium  permanganate,  in  which  the  permanganic 
acid  has  been  set  free  by  an  acid:  KMn04  +  H2S04  <=±  HMn04  + 
KHS04,  is  rapidly  reduced.  The  permanganic  acid,  with  excess 
of  sulphuric  acid,  tends  to  undergo  the  first  of  the  following  changes, 
provided  a  substance,  such  as  hydrogen  peroxide,  is  present  which 
can  take  possession  of  the  oxygen  that  would  remain  as  a  balance: 

2HMn04  +  2H2S04  -*  2MnS04  +  3H20  (+  50).  (1) 
_  (5O)  +  5H2O2  ->  5H2O  +  502.  _  (2) 
2HMnO4  +  2H2S04  +  5H202  ->  2MnSO4  +  8H20  +  502. 

Exercises.  —  1.  What  volume  of  ozone  will  be  taken  up  by  100 
c.c.  of  water  at  12°  from  a  stream  of  oxygen  containing  7.5  per  cent 
of  ozone  (p.  129)? 

2.  At  what  temperature  will  a  ten  per  cent  aqueous  solution  of 
hydrogen  peroxide  freeze  (p.  214)? 

3.  Write  the  thermochemical  equations  for  oxidation  of  indigo 
by  hydrogen  peroxide  (pp.  221,  224). 

4.  How  many  times  its  own  volume  of  oxygen  gas  will  a  3  per 
cent  solution  of  hydrogen  peroxide  give  off  when  treated  with: 
(a)  platinum  powder  (p.  223);    (6)  sulphuric  acid  and  potassium 
permanganate? 


CHAPTER  XVIII 
IONIZATION 

Introductory.  —  As  we  have  seen,  acids,  bases,  and  salts,  when 
dissolved  in  water,  interact  with  one  another  by  interchanging 
radicals  (p.  148).  We  have  also  learned  that  the  same  solutions 
have  abnormal  values  for  their  freezing-points  and  for  two  other 
properties.  These  facts  indicate  dissociation  into  the  radicals  (p. 
216).  Now  precisely  these  solutions  have  a  property  which  is  not 
shared  by  any  other  solutions,  namely,  that  of  being  conductors  of 
electricity  and  suffering  chemical  decomposition  by  the  passage  of  the 
current.  Such  solutions  are  called,  in  consequence,  electrolytes, 
and  the  process  is  named  electrolysis.  Now  the  natural  inference 
from  the  foregoing  facts  is  that  the  electricity  is  carried  by  the 
liberated  radicals.  Our  first  aim  in  the  present  chapter  is  to  show, 
by  a  study  of  the  chemical  changes  taking  place  in  electrolysis,  that 
this  inference  is  correct.  We  then  proceed  to  discuss  the  nature  of 
ions  as  a  kind  of  molecules.  Next,  we  devote  ourselves  to  the 
explanation  of  electrolysis,  to  the  equilibrium  between  the  ions  and 
the  remaining,  undissociated  molecules,  and  to  conductivity  phe- 
nomena as  a  means  of  measuring  the  fraction  ionized.  Finally,  we 
deduce  the  relation  between  extent  of  ionization  and  chemical 
activity. 

Incidentally,  the  facts  to  be  given  provide  the  means  of  under- 
standing the  electrolytic  processes,  many  of  them  of  great  impor- 
tance in  chemical  industries,  to  which  frequent  reference  is  made  in 
later  chapters. 

Non- Electrolytes.  —  To  clear  the  ground,  we  should  first  note 
the  fact  that  only  solutions  (as  a  rule)  possess  both  of  the  properties 
in  question,  namely  that  of  conducting  and  that  of  being  decom- 
posed by  the  current.  Some  substances,  notably  the  metals  and 
materials  like  carbon,  are  conductors.  But  they  are  not  changed 
chemically  by  the  current.  Again,  single  substances,  even  when 
they  are  such  as,  if  mixed,  yield  electrolytes,  are  not  conductors  at 

226 


IONIZATION  227 

ordinary  temperatures.  Thus  hydrogen  chloride,  whether  gaseous 
or  liquefied,  is  a  nonconductor,  and  water  is  a  very  feeble  conduc- 
tor, although  the  solution  of  the  two  conducts  exceedingly  well. 
Dry  acids,  bases,  and  salts,  except  when  at  a  high  temperature  and 
fused,  are  likewise  nonconductors.  Furthermore,  even  amongst 
solutions,  not  all  are  conductors.  Solutions  of  sugar  and  other 
substances  of  the  same  class  (p.  213),  which  have  normal  freezing- 
points,  are  nonconductors.  Only  solutions  of  acids,  bases,  and 
salts  in  certain  specified  solvents,  of  which  the  commonest  is  water, 
are  electrolytes  at  ordinary  temperatures. 

Chemical  Changes  Taking  Place  in  Electrolysis:  at  the 
Electrodes.  —  When  the  wires  from  a  battery  are  attached  to 
platinum  plates  immersed  in  any  electrolyte  (e.g.,  Fig.  65,  p.  155),  we 
observe  that  the  products  appearing  at  the  two  electrodes  are 
always  different.  They  may  be  of  several  kinds  physically,  and 
will  be  secured  for  examination  variously  according  to  their  nature. 
Thus,  when  they  are  gases,  which  are  not  too  soluble,  they  may  be 
collected  in  inverted  tubes  filled  with  the  solution.  Solids,  if  in- 
soluble in  the  liquid,  will  either  remain  attached  to  the  electrode  or 
fall  to  the  bottom  of  the  vessel  as  precipitates.  Soluble  substances, 
on  the  other  hand,  will  usually  not  be  visible.  They  may  be 
handled  by  interposing  a  porous  partition  of  some  description 
which  will  restrain  the  diffusion  of  the  dissolved  body  away  from 
the  neighborhood  of  the  electrode,  while  not  interfering  appreciably 
with  the  passage  of  the  current.  Surrounding  one  electrode  with 
a  porous  battery  jar  is  a  convenient  method  for  effecting  this. 

Of  the  various  illustrations  which  we  have  encountered,  the  elec- 
trolysis of  hydrochloric  acid  (p.  155)  happens  to  have  been  the  only 
one  which  delivered  both  components  of  the  solute  with  a  minimum 
of  modification  at  the  electrodes: 

Neg.  Wire,  H2< H.C1 >C12,  Pos.  Wire. 

Hydrogen  does  not  interact  with  water,  and  chlorine  interacts  very 
incompletely,  so  that  the  molecular  substances  H2  and  C12  are 
promptly  formed  from  the  elements  H  and  Cl  which  are  liberated. 
The  chlorides,  bromides,  and  iodides  of  those  metals  which  do  not 
interact  with  water  (p.  60)  give  equally  simple  results: 
Neg.  Wire,  Cu< Cu.Br2 >Br2,  Pos.  Wire. 


228  COLLEGE    CHEMISTRY 

Thus  the  solute  seems  to  be  split  into  its  radicals  and,  in  elec- 
trolysis, the  radicals,  if  they  do  not  interact  with  water,  are  set 
free.  A  substance  thus  set  free  is  called  a  primary  product  of 
the  electrolysis.  In  the  foregoing  instances  both  products  are 
primary. 

Usually  the  chemical  change  is  more  complex.  Thus,  when 
dilute  sulphuric  acid  is  electrolyzed,  hydrogen  and  oxygen  are 
liberated  at  the  negative  and  positive  electrodes,  respectively. 
But  these  products  do  not  account  for  the  whole  of  the  constitu- 
ents (H2S04).  We  therefore  proceed  to  examine  the  materials  in 
solution  round  the  electrodes.  It  is  found  that,  as  the  action 
progresses,  sulphuric  acid  accumulates  round  the  positive  wire, 
while  the  liquid  in  the  neighborhood  of  the  other  pole  is  gradually 
depleted  of  this  substance.  In  view  of  this  fact  we  easily  explain 
the  phenomenon.  Evidently  the  substance  divides  into  its  radi- 
cals, H  and  864,  but  864,  not  being  a  known  substance,  must 
interact  with  the  water  to  produce  sulphuric  acid  and  oxygen: 
2SO4  +  2H2O  ->  2H2SO4  +  02.  The  whole  change  may  therefore 
be  tabulated  as  follows: 

Neg.  Wire,  H2< H2.S04 >02  andH2S04,  Pos.  Wire. 

Hence  the  hydrogen  is  a  primary  product,  but  the  oxygen  and  sul- 
phuric acid  are  secondary  products.  All  acids  give  hydrogen  alone 
at  the  negative  electrode,  whatever  may  be  the  product  at  the 
positive. 

If  we  electrolyze  cupric  nitrate  solution,  we  obtain  a  red  deposit 
of  metallic  copper  on  the  negative  plate  and  at  the  positive  end 
oxygen  and  nitric  acid  are  formed.  We  infer,  therefore,  that  the 
division  of  the  original  molecule  was  into  Cu  and  N03,  but  that  the 
latter  interacted  with  the  water :  4NO3  +  2H20 >  4HN03  +  02 : 

Neg.  Wire,  Cu< Cu.(N03)2 »02  and  HN03,  Pos.  Wire. 

With  a  solution  of  potassium  chloride  we  find  hydrogen  and 
chlorine  appearing  at  the  negative  and  positive  electrodes,  re- 
spectively. Litmus  paper,  however,  shows  the  presence  in  the 
solution  of  a  base  (potassium  hydroxide,  KOH)  at  the  negative 
end.  We  infer  that  the  parts  of  the  parent  molecules  are  K  and 
Cl.  The  former,  since  it  resembles  sodium  in  being  much  more 
active  than  hydrogen  (p.  60),  is  more  difficult  to  liberate.  Hence 


IONIZATION  229 

hydrogen  is  liberated  instead,  and  potassium  hydroxide  remains 
in  the  liquid:  2K  +  2HOH  ->  2KOH  +  H2: 

Neg.  Wire,  H2  and  KOH< K.C1 >C12,  Pos.  Wire. 

We  are  confirmed  in  this  explanation  when  we  employ  a  solution 
containing  a  mixture  of  salts  of  copper  and  silver.  The  latter, 
being  the  less  active  metal,  is  first  deposited,  alone.  The  copper 
is  liberated  only  after  all  the  silver  has  been  set  free. 

Having  now  before  us  the  results  of  electrolyzing  some  typical 
substances,  we  bring  these  results  into  relation  with  the  facts 
described  in  Chapter  XVI.  Acids  contain  hydrogen  which  pos- 
sesses certain  specific  properties  (p.  210),  and  by  electrolysis  all 
acids  divide  so  as  to  give  up  this  constituent  alone  at  one  electrode. 
The  evidence  that  the  other  radical  has  different  electrical  proper- 
ties which  carry  it  to  the  opposite  plate  is  conclusive.  Again,  salts 
undergo  double  decomposition  in  which  they  exchange  radicals 
with  acids,  bases,  and  other  salts  (p.  211),  and  we  find  that  it  is 
these  very  radicals  which  are  withdrawn  from  the  solution  by  the 
influence  of  the  electricity.  Furthermore,  the  radicals  exist  free  in 
the  solution,  being  formed  by  dissociation  of  the  molecules  (p.  216). 
Hence  the  function  of  the  electricity  seems  simply  to  consist  in  sifting 
apart  the  two  kinds  of  free  radicals  which  each  solution  contains. 
It  only  remains  for  us  to  explain  in  detail  the  sifting  action  of  the 
current.  Before  turning  to  the  explanation  of  this  phenomenon, 
however,  there  is  one  question  which  may  be  answered  in  passing. 
Since  a  solution  may  eventually  be  cleared  of  all  the  hydrochloric 
acid,  for  example,  which  it  contains,  we  should  like  to  know  how 
the  free  radicals  in  the  center  of  the  cell  reach  the  electrodes. 

Ionic  Migration.  —  To  know  how  the  free  radicals  reach  the 
electrodes,  all  that  is  necessary  is  to  take  a  material,  one  (or  both) 
of  whose  radicals  is  a  colored  substance,  and  watch  the  move- 
ment of  the  colored  material  as  it  drifts  towards  the  electrode. 
Most  salts  which  give  colored  solutions  are  suitable.  In  very 
dilute  cupric  sulphate  solution,  for  example,  a  freezing-point 
determination  shows  that  the  depression  has  practically  double 
the  normal  value.  In  other  words,  the  dissociation  into  the 
radicals,  CuS04^.(Cu)  +  (S04),  is  almost  complete.  Now,  the 
blue  color  of  the  solution  cannot  be  due  to  the  few  remaining 


230 


COLLEGE    CHEMISTRY 


molecules  of  CuSO4,  for  anhydrous  cupric  sulphate  is  colorless. 
Nor  is  it  due  to  the  color  of  the  (S04)  radicals,  for  dilute  potassium 
sulphate  and  dilute  sulphuric  acid  are  both  colorless.  On  the  other 
hand,  all  cupric  salts,  in  dilute  solution,  have  the  same  tint.  The 
color  is  therefore  that  of  the  free  cupric  radical  (Cu).  In  order 
most  clearly  to  see  the  motion  of  the  cupric  radical,  we  place  the 
cupric  sulphate  solution  in  the  middle  of  the  space  between  the 
electrodes,  and  place  between  it  and  the  latter  a  colorless  con- 
ducting solution.  The  motion  of  the  blue  material  across  the 
boundary  may  then  be  easily  observed. 

The  most  convenient  arrangement  is  to  dissolve  the  cupric  sul- 
phate in  warm  water  containing  about  5  per  cent  of  agar-agar  (a 
gelatine  obtained  in  China  from  certain  sea-weeds),  and  to  fill  with 
this  mixture  the  lower  part  of  a  U-tube  (Fig.  79).  The  setting  of 
the  jelly  prevents  subsequent  mixing  of  the  cupric  sulphate  system 

of  materials  with  the  rest  of  the 
filling  of  the  tube,  and  the  conse- 
quent disappearance  of  the  bound- 
ary. A  few  grains  of  charcoal  are 
scattered  on  the  surface  of  the 
jelly  to  mark  the  present  limits  of 
the  colored  substance,  and  a  solu- 
tion of  some  colorless  electrolyte, 
such  as  potassium  nitrate,  is 
added  on  each  side.  To  prevent 
agitation  of  the  liquid  by  the 
effervescence  at  the  electrodes,  it 
is  well  to  use  agar-agar  with  the 
lower  part  of  the  colorless  liquid 
also.  The  whole  is  finally  placed 
in  ice  and  water,  to  prevent  melt- 
ing of  the  jelly  by  the  heat  caused  by  resistance,  and  the  current 
is  then  turned  on. 

After  a  time,  we  observe  that  the  blue  cupric  radicals  ascend 
above  the  mark  on  the  negative  and  descend  away  from  it  on  the 
positive  side.  In  each  case  there  is  no  shading  off  in  the  tint.  The 
motion  of  the  whole  aggregate  of  colored  radicals  occurs  in  such  a 
way  that,  if  the  contents  of  the  tube  were  not  held  in  place  by  the 
jelly,  we  should  believe  that  a  gradual  motion  of  the  entire  blue 


FIG.  79. 


IONIZATION  231 

solution  was  being  observed.  With  a  current  of  110  volts,  and  a 
16-candle-power  (one-half  ampere)  lamp  in  series  with  the  cell,  the 
effect  becomes  apparent  in  a  few  minutes. 

Although  the  (S04)  radicals  are  invisible,  we  may  safely  infer 
that  they  are  drifting  towards  the  positive  electrode.  Indeed,  this 
can  be  demonstrated  by  interposing  a  shallow  layer  of  jelly  con- 
taining some  barium  salt  a  little  distance  above  the  charcoal  layer 
on  the  positive  side.  When  the  (864)  reaches  this,  barium  sul- 
phate BaSO4  begins  to  be  precipitated  and  the  layer  becomes 
cloudy.  In  similar  ways  the  progress  of  other  colorless  ions  may 
be  rendered  visible. 

It  appears,  therefore,  that  electrolysis  is  not  a  local  phenomenon, 
going  on  round  the  electrodes  only,  but  that  the  whole  of  the 
products  of  the  dissociation  of  the  solute  are  set  in  motion.  It  is  on 
account  of  this  remarkable  property  of  traveling  or  migrating 
towards  one  or  other  of  the  electrodes  that  the  individual  atoms 
(like  Cu),  or  groups  of  atoms  (like  SO4),  have  been  named  ions 
(Gk.,  going).  The  term  was  first  applied  by  Faraday  to  the 
materials  liberated  round  the  electrodes. 

Different  ionic  substances  move  with  different  speeds  when  pro- 
pelled by  the  same  current.  The  hydrogen  radical  of  acids  (H)  is 
the  most  speedy,  the  hydroxyl  radical  of  bases  (OH)  comes  next. 
These  are,  respectively,  about  five  and  two  and  one-half  tunes  as 
fast  as  any  other  ions.  The  actual  speeds  of  several  ions,  in  dilute 
solutions  at  18°,  when  driven  by  a  potential  difference  of  1  volt 
between  plates  1  cm.  apart,  expressed  in  cm.  per  hour  is:  H  10.8, 
OH  5.6,  Cu  1.6,  S04  1.6,  K  2.05,  Cl  2.12. 

The  Nature  of  Ions:  Faraday's  Law.  —  That  the  molecules 
of  certain  classes  of  substances,  although  seemingly  without  chemi- 
cal interaction  with  the  water  in  which  they  are  dissolved,  should 
nevertheless  be  decomposed  by  the  influence  of  the  water,  is 
strange,  but  not  inconceivable.  Heating  produces  a  somewhat 
similar  effect  on  many  substances.  The  novel  fact,  for  which  an 
explanation  is  demanded,  is  that  the  molecules  of  the  products  of 
the  dissociation  appear  to  be  attracted  by  electrically  charged 
plates,  which  have  been  lowered  into  the  solution,  while  molecules 
of  dissolved  sugar,  for  example,  are  not  so  attracted.  Now  the  only 
bodies  which  we  find  to  be  conspicuously  attracted  by  electrically 


232  COLLEGE    CHEMISTRY 

charged  objects  are  bodies  which  are  already  provided  with  electric 
charges  of  their  own.  Thus  we  are  led  to  add  the  idea  that  sub- 
stances which  undergo  dissociation  in  solution  divide  themselves 
into  a  special  kind  of  electrically  charged  molecules. 

Since  the  solution,  as  a  whole,  has  itself  no  charge,  equal  quan- 
tities of  positive  and  negative  electricity  must  be  produced: 

HC1  «=*  H+  +  OP  NaCl  ^  Na+  +  CP  NaOH  <=»  Na+  +  OH". 

This  means  that  bivalent  radicals,  on  dissociation,  will  become  ions 
carrying  a  double  charge  and  trivalent  ions  must  carry  a  triple 
charge : 

CuCl2  ?=»  Cu++  +  2C1~        CuS04  ?±  Cu++  +  S04= 
K2SO4  ?±  2K+  +  S04=         FeCl3  <=*  Fe+++  +  3C1~ 

In  these  equations,  the  coefficients  multiply  the  charges  as  well  as 
the  radicals  bearing  the  charges,  and  it  will  be  seen  that  the  num- 
bers of  +  and  —  charges  produced  by  each  dissociation  are  equal. 
Hence,  univalent  ions  all  possess  equal  quantities  of  electricity,  and 
other  ions  bear  quantities  greater  than  this  in  proportion  to  their 
valence.  This  is  an  inevitable  inference  from  the  electrical  neu- 
trality of  all  solutions.  An  ion  is  therefore  an  atom  or  group  of 
atoms  bearing  an  electric  charge. 

This  conclusion  is  confirmed  by  actual  measurement.  When 
hydrochloric  acid  is  electrolyzed,  35.46  g.  (=  Cl)  of  chlorine  are 
liberated  for  every  1.008  g.  (=  H)  of  hydrogen.  But  when  cupric 
chloride  CuCl2  is  substituted,  for  every  35.46  g.  ( =  Cl)  of  chlorine 
set  free,  only  31.78  g.  (=  \  Cu  =  \  63.57)  of  copper  is  deposited. 
The  law,  discovered  by  Faraday,  is  that :  equal  quantities  of  electric- 
ity liberate  equivalent  quantities  of  the  ions  (equivalent,  p.  65, 
not  atomic  or  molecular). 

To  show  that  this  view  of  the  nature  of  the  ions  is  adequate,  we 
next  apply  it  to  the  explanation  of  the  phenomena  of  electrolysis. 
After  that  some  seeming  objections  will  be  discussed. 

Application  to  the  Explanation  of  Electrolysis.  —  A  bat- 
tery is  a  machine  which  maintains  two  points,  its  poles,  or  two 
wires  connected  with  them,  at  a  constant  difference  of  potential. 
One  cell  of  a  lead  storage  battery,  for  example,  maintains  a  poten- 
tial difference  of  about  two  volts.  When  the  wires  are  joined, 


IONIZATION  233 

directly  or  indirectly,  the  poles  are  immediately  discharged,  but  the 
:ell  continuously  reproduces  the  difference  in  potential  by  generat- 
ing fresh  electricity.  Now  the  effect  of  immersing  two  plates,  one 
of  which  is  kept  by  the  battery  at  a  definite  positive  potential  and 
the  other  at  a  definite  negative  potential,  into  a  liquid  filled  with 
multitudes  of  minute,  suspended  bodies,  already  highly  charged, 
may  easily  be  foreseen. 

The  figure  (Fig.  80)  will  convey  some  idea  of  the  behavior  of  the 
parts  of  a  system  such  as  we  have  imagined.  The  electrodes  are 
marked  —  and  +.  The  negatively  charged  plate  attracts  all  the 
positively  charged  particles  Cathode  + 

in  the  vessel  and,  although         ._        <— cation  =  Ag 
these  particles  were  in  con-       ^*""^  anion  =  NO3— » 

tinuous  and  irregular  mo- 
tion, they  at  once  begin  to 
drift  toward  the  plate  in 
question.  On  the  other 
hand,  the  negatively 


§ 

O 


LiCKllJ.\JLj  UJLJLV^  J-iV^^CAIUi  J    J.    J  V [Illlll  ^S 

charged    particles    are    re- 
pelled  by   this   plate    and  FIG.  so. 
attracted   by   the   positive 

plate,  so  that  they  drift  in  the  opposite  direction.  Those  which  are 
nearest  each  plate,  on  coming  in  contact  with  it,  will  have  their 
charges  of  electricity  neutralized  by  the  opposite  charge  on  the  plate, 
turning  thereby  into  the  ordinary  free  forms  of  the  matter  of  which 
they  are  composed.  The  continuous  removal  of  the  electrical  charges 
of  the  plates  through  contact  with  ions  of  the  opposite  charge  fur- 
nishes occasion  for  recharging  of  the  plate  from  the  battery,  and  thus 
gives  rise  to  a  continuous  current  in  each  wire.  Again,  the  continu- 
ous drifting  of  positively  and  negatively  charged  particles  in  oppo- 
site directions  through  the  liquid,  constitutes  what,  in  the  view  of 
all  external  means  of  observation,  appears  to  be  an  electical  current 
in  the  liquid  also.  A  magnetized  needle,  for  example,  which  is  de- 
flected when  brought  near  to  one  of  the  wires  of  the  battery,  is  in- 
fluenced in  the  same  way  by  being  brought  over  the  liquid  between 
the  electrodes.  The  illusion,  so  to  speak,  of  an  electric  current  i 
complete,  although  in  reality  it  is  a  convection  of  electricity  that  ] 
taking  place.  Furthermore,  the  quantity  of  electricity  being  trans- 
ported across  any  section  of  the  whole  system  is  the  same  as  that 


234  COLLEGE    CHEMISTRY 

across  any  other,  whether  this  section  be  taken  through  one  of  the 
wires,  through  the  electrolyte,  or  even  through  the  battery  at  any 
point.  As  fast  as  the  ions  are  thus  annihilated  as  such,  the  undis- 
sociated  molecules  (mingled  with  the  ions,  but  not  shown  in  the  fig- 
ure) dissociate  and  produce  fresh  ones,  as  in  all  chemical  equilibria. 
Eventually,  by  continuing  the  process  long  enough,  if  the  substances 
set  free  are  actually  deposited  and  do  not  go  into  solution  again 
in  any  form,  the  liquid  can  be  entirely  deprived  of  the  whole  of  the 
solute  which  it  contains. 

The  analogy  to  the  transportation  of  a  fluid  like  water  is  notice- 
able, although  not  complete.  Water  may  be  transported  in  three 
ways.  It  may  flow  through  a  pipe,  it  may  pass  by  pouring  freely 
from  one  container  to  another,  and  it  may  be  carried  in  vessels. 
Thus  a  stream  of  water,  essentially  continuous,  might  be  arranged, 
in  which  part  of  the  passage  took  place  through  the  pipes,  part  by 
pouring  from  the  pipes  into  buckets,  and  part  by  the  carrying  of 
those  buckets  between  the  ends  of  the  pipes.  The  quantity  of 
water  passing  a  given  point  per  minute  in  this  system  would  be  the 
same  at  every  part,  although  the  actual  method  by  which  the  water 
was  transported  past  the  various  points  might  be  different.  In 
such  a  disjointed  circuit  we  suppose  the  electricity  to  move  when 
carried  from  a  battery  through  an  electrolytic  cell.  It  flows  in 
the  wire,  passes  by  discharge  between  the  pole  and  the  ion,  and  is 
transported  upon  the  ions  in  the  liquid.  The  parallel  is  imperfect, 
however,  because  we  have  used  the  conception  of  two  electric  fluids 
and  because  the  ions  are  already  charged  in  the  solution,  and  before 
any  connection  with  the  battery  is  made.  They  do  not,  so  to  speak, 
transport  the  electricity  of  the  battery,  but  their  own. 

Questions  Suggested  by  this  Explanation.  —  1.  The  ques- 
tion was  raised  (p.  217),  as  to  how  we  can  imagine  separate  atoms 
of  sodium  to  exist  in  water  without  acting  upon  it,  as  the  metal 
sodium  usually  does.  But  the  ions  of  sodium  in  sodium  chloride 
solution  are  not  metallic  sodium.  They  bear  large  charges  of 
electricity.  They  possess  an  entirely  different,  and  in  fact,  by 
measurement,  much  smaller  amount  of  chemical  energy  than  free 
sodium.  And,  as  we  have  seen,  the  properties  of  a  substance  are 
determined  as  much  by  the  energy  it  contains  as  by  the  kind  of 
matter.  Metallic  sodium  and  ionic  sodium  are,  simply,  different 


IONIZATION  235 

substances.  Besides,  when  metallic  sodium  acts  on  water,  it  turns 
into  the  ionic  sodium  of  sodium  hydroxide  (Na+  -f  OH~^±  NaOH). 
Ionic  sodium  (Na+)  from  sodium  chloride  is,  therefore,  already  in 
the  very  state  which  metallic  sodium  reaches  by  interaction  with 
water,  and  is  in  no  need  of  trying  to  enter  that  state. 

2.  We  think  of  hydrogen  chloride  and  common  salt  as  exceed- 
ingly stable  substances,  and  are  averse  to  believing  that  precisely 
these  compounds  should  be  highly  dissociated  by  mere  solution  in 
water.     But  it  must  be  remembered  that  in  solution  they  undergo 
chemical  change  very  easily,  and  it  is  only  in  the  dry  form  that 
they  show  unusual  stability. 

3.  Again,  why  do  not  the  ions  combine,  in  response  to  the  at- 
tractions of  their  charges?     The  answer  is  that  they  do  combine, 
but  the  rate  at  which  combination  takes  place  is  no  greater  than 
that  at  which  the  molecules  decompose,  so  that  on  the  whole  the 
proportion  of  ions  to  molecules  remains  unchanged. 

4.  It  might  appear  that  the  idea  that  bodies  could  retain  high 
charges  in  the  midst  of  water  is  contrary  to  all  experience.     It 
must  be  remembered,  however,  that  the  molecular,  pure  water, 
which  separates  the  ions  from  one  another,  is  a  perfect  nonconduc- 
tor.    The  moisture  which  covers  electrical  apparatus  and  causes 
leakage  of  static  electricity  is  not  pure  water,  but  a  dilute  solution, 
containing  carbonic  acid  (p.  91)  and  materials  from  the  glass  of 
which  the  apparatus  is  made  (p.  92).    It  conducts  away  the 
charge  electrolytically,  by  means  of  the  ions  it  contains,  and  not 
by  itself  acting  as  a  conductor. 

5.  Finally,  when  we  dissolve  an  electrically  neutral  salt  in  water, 
whence  do  the  radicals  obtain  the  electric  charges?    We  now  know 
that  an  atom,  say,  of  sodium,  contains  a  minute  nucleus  of  positive 
electricity,  which  contains  most  of  the  mass  of  the  atom.    Outside 
of  this  nucleus,  there  are  particles  of  negative  electricity,  called 
electrons   (q.v.)}  each    having  a  mass   about  one-eighteen  hun- 
dredth (T*W)  of  that  of  an  atom  of  hydrogen.    An  ion  of  chlorine 
(Cl~)  consists,  therefore,  of  an  atom  of  chlorine  plus  one  electron 
(Cl  +  c).     An  ion  of  sodium  is  an  atom  of  sodium  minus  one 
electron  (Na  -  e)  and  has  thus  an  excess  of  one  unit  positive 
charge  in  the  nucleus.    When  these  two  ions  combine,  the  result- 
ing molecule  NaCl  is  neutral. 


236 


COLLEGE    CHEMISTRY 


Resume  and  Nomenclature.  —  The  dissociation  of  molecules 
into  ions  is  named  ionization.  The  substances  of  the  three  classes 
which  alone  are  ionized  may  be  designated  ionogens.  An  ion  may 
be  defined  as,  a  molecule  bearing  negative  or  positive  charges  of 
electricity  in  proportion  to  its  valence,  and  formed  through  the 
dissociation  of  an  ionogen  by  a  solvent  like  water. 

Each  molecule  of  the  solute  gives  two  kinds  of  ions  with  opposite 
charges.  These  two  are  forthwith  distinct  and  independent  sub- 
stances, save  that  the  attractions  of  the  charges  prevent  any  con- 
siderable separation  by  diffusion.  They  differ  from  non-ionic 
substances  of  the  same  material  composition  when  such  are  known. 
The  electrical  charge  is  one  of  the  essential  constituents  and,  when 
it  is  removed,  the  properties  alter  entirely.  Thus  we  have  two 
kinds  of  hydrogen,  gaseous  molecular  hydrogen  (H2),  and  ionic 
hydrogen  (H+),  with  entirely  different  chemical  properties. 

The  radicals  and  their  chemical  behavior  are  real,  and  all  the 
peculiarities  of  aqueous  solutions  of  acids,  bases,  and  salts  are 
experimental  facts.  We  now  have  experimental  knowledge  of 
the  minute  parts  of  bodies.  Molecules  are  units  which  are  not 
commonly  disintegrated  by  vaporization  (p.  102);  ions,  those 
which  are  not  commonly  disintegrated  in  double  decomposition  in 
solution;  atoms,  those  which  are  not  commonly  disintegrated  in 
any  chemical  action.  The  ionic  explanation  was  first  suggested 
as  an  hypothesis  by  Svante  Arrhenius,  a  Swedish  chemist,  in 
1887. 

Since  ionic  hydrogen,  ionic  chlorine,  etc.,  are  entirely  different  in 
physical  and  chemical  properties  from  the  corresponding  free  ele- 
ments, they  should  receive  separate  names.  When  it  is  incon- 
venient to  say  "ionic  hydrogen,"  "ionic  nitrate  radical"  (N03~), 
etc.,  the  following  names  will  be  used  for  the  ionic  substances: 


Sym- 
bol. 

Name  of 
Substance 

Anion  of 

Symbol. 

Nameof 
Substance. 

Cation  of  Salts  of 

S04= 

Sulphate-ion 

Sulphates 

Na+ 

Sodium-ion 

Sodium 

cr 

Chloride-ion 

Chlorides 

Fe-H+ 

Ferric-ion 

Ferric  iron 

Hsor 

Hydrosulphate-ion 

Biaulphates 

NH4+ 

Ammonium-ion 

Ammonium 

OH~ 

Hydroxide-ion 

Hydroxides 

Fe++ 

Ferrous-ion 

Ferrous  iron 

(bases) 

H+ 

Hydrogen-ion 

Hydrogen  (acids) 

IONIZATION  237 

In  using  these  terms,  note  that  sodium-ion  (with  the  hyphen)  is  the 
name  of  the  substance,  and  not  of  the  charged  atom.  When 
speaking  in  terms  of  ions  as  particles,  therefore,  we  may  not  say 
"a  sodium-ion,"  any  more  than  we  should  say  "an  ionic  sodium" 
or  "ionic  sodiums."  To  describe  the  charged  molecule,  we  must 
write  "a  sodium  ion,"  "sodium  ions,"  "chlorate  ions,"  etc. 

Faraday  distinguished  the  two  kinds  of  material  which  proceed 
with  and  against  the  positive  current  by  name.  His  terminology  is 
still  used.  Ions  which  proceed  in  the  same  direction  as  the  positive 
current  (Fig.  80)  are  called  cations  (Gk.,  down).  Such  are  H+, 
Cu++,  K+,  NH4+.  They  are  metallic  elements,  or  groups  which  play 
the  part  of  a  metal.  The  electrode  (Gk.,  a  path  for  electricity)  upon 
which  they  are  deposited,  the  negative  electrode,  is  spoken  of  as 
the  cathode  (Gk.,  the  way  down). 

The  particles  which  move  in  the  direction  of  the  negative  current, 
and  against  that  of  the  positive,  are  named  anions  (Gk.,  up).  The 
ions  Cl~,  N03~,  SO4=,  Mn04~  are  of  this  kind.  They  are  usually 
composed  of  non-metals,  although  sometimes,  as  in  MnCX",  the  con- 
stituents may  be  partially  metallic.  They  are  set  free  at  the  posi- 
tive electrode,  which  is  therefore  named  the  anode  (Gk.,  the  way 
up).  Chemists  speak  of  metals  and  non-metals  as  positive  and 
negative  elements,  respectively,  even  when  electrical  relations  are 
not  directly  in  question,  and  ions  are  not  concerned. 

Actual  Quantities  of  Electricity  Concerned.  —  The  units 
of  electrical  energy  are  the  coulomb,  which  is  the  unit  of  quantity, 
and  the  volt,  which  is  the  unit  of  difference  of  potential  (or  pressure, 
so  to  speak).  Faraday's  law  has  to  do  only  with  the  former. 
Equal  numbers  of  coulombs  liberate  equivalent  weights  of  all 
elements,  but  different  voltages  are  required  to  decompose  differ- 
ent compounds,  according  to  then-  stability  (see  Chap.  XXXIX). 

To  liberate  1.008  g.  of  hydrogen,  or  one  equivalent  of  any  other 
element,  96,540  coulombs  of  electricity  are  needed.  The  charges 
on  1.008  g.  of  hydrogen  ions  must,  therefore,  equal  this  amount. 
There  are  6.07  X  1023  molecules  of  hydrogen  in  22.4  liters  (H2)  and 
therefore  in  2.016  g.  of  the  gas.  A  simple  calculation  shows  there- 
fore that  each  coulomb  is  distributed  over  about  63  X  1017  ions 
of  hydrogen. 

A  current  of  1  coulomb  per  second  is  called  1  ampere.     Thus, 


238  COLLEGE    CHEMISTRY 

the  current  passing  through  a  1-amp.  lamp  (or  2  half-ampere  16- 
c.p.  lamps  in  parallel)  will  liberate  1.008  g.  (11.2  liters)  of  hydro- 
gen in  96,540  seconds,  or  26  hours  and  49  minutes.  The  same 
current  will  liberate  107.88  g.  of  silver  (Ag1),  or  31.78  g.  of  copper 
(Cun/2)  from  cupric  sulphate  in  the  same  time.  A  current  of 
5  amperes  will  accomplish  the  same  result  in  one-fifth  of  the  time. 

Applications:  Ionic  Equilibrium.  —  Since  the  ions  are  chemi- 
cally different  from  their  parent  molecules,  their  formation  repre- 
sents a  variety  of  chemical  change.  The  change  does  not  involve 
any  chemical  interaction  with  the  water.  It  is  simply  a  dissocia- 
tion, i.e.,  reversible  decomposition  of  the  dissolved  substance. 

From  the  fact  that  the  proportion  of  molecules  ionized  is  shown 
to  become  greater  as  more  and  more  of  the  solvent  is  added  (p.  215), 
and  that  removal  of  the  solvent  diminishes  the  proportion  of  ions  to 
molecules,  and  finally  leaves  the  substance  entirely  restored  to 
the  molecular  condition,  we  know  that  this  is  a  reversible  action  and 
therefore  a  true  dissociation.  The  molecules  and  their  ions  adjust 
themselves  like  the  constituents  in  any  case  of  chemical  equilibrium 
(pp.  177-182): 

NaCl«=*Na+  +  CT. 

The  chemical  behavior  of  substances  in  ionic  equilibrium  will  be 
discussed  in  the  next  chapter  (see  p.  249). 

*  The  mode  of  formulation  previously  used  (p.  183)  may  be 
employed  here.  If  [NaCl],  [Na+],  and  [Cl~]  stand  for  the  molec- 
ular concentrations  (numbers  of  moles  per  liter)  at  equilibrium  of 
the  molecules,  and  the  two  ions,  respectively,  we  have  an  equilib- 
rium constant  (cf.  p.  184),  in  this  case  called  the  ionization 
constant  : 


[NaCl] 

When  we  dissolve  a  single  substance  which  gives  only  two  ions,  the 
molecular  concentrations  of  the  ions  are  necessarily  equal.  When 
some  other  ionogen  with  a  common  ion  is  present,  however,  the 
values  of  [Na+]  and  [Cl~]  will  be  different. 

*  The  content  of  this  paragraph  is  referred  to  in  Chap.  XX,  but  is  not  em- 
ployed systematically  until  Chap.  XXXV  is  reached. 


IONIZATION 


239 


Applications:  To  the  Interpretation  of  Conductivity 
Measurements.  —  We  have  seen  that  when  the  solution  of  an 
ionogen  is  diluted,  the  proportion  of  ions  to  undissociated  molecules 
increases,  while  removal  of  a  part  of  the  solvent  has  the  opposite 
effect  (p.  215).  Now,  a  change  in  the  number  of  ions  naturally 
modifies  the  capacity  of  the  liquid  for  carrying  electricity,  so  that 
observation  of  the  changes  in  the  conductivity  of  a  solution,  when 
the  concentration  is  altered,  supplies  the  simplest  means  of  studying 
the  phenomena  of  ionization. 

A  glass  trough  and  amperemeter  *  (Fig.  81)  may  be  used  to  illus- 
trate this  principle.  The  electrodes  are  long  strips  of  copper  foil, 
which  pass  down  at  the  ends  of  the  trough.  After  placing  the  two 


Fia.  81. 

instruments  in  circuit  with  a  source  of  electricity,  we  first  pour  very 
pure  water  into  the  cell.  With  this  arrangement,  the  ampere- 
meter does  not  indicate  the  passage  of  any  current  of  electricity. 
Concentrated  (36  per  cent)  hydrochloric  acid  is  next  cautiously 
added  through  a  long-stemmed  dropping  funnel,  so  that  it  forms  a 
shallow  layer  below  the  water,  and  the  funnel  is  withdrawn.  The 
situation  at  this  stage  is  that  a  definite  amount  of  hydrogen 
chloride  dissolved  in  a  small  amount  of  water  fills  what  was  before 
a  nonconducting  gap  in  the  electric  circuit.  The  deflection  of  the 
needle  in  the  amperemeter  indicates  that  a  certain  current  of 
electricity  is  able  to  pass  through  this  acid.  When  we  now  stir  the 
*  An  amperemeter  of  low  resistance,  0.5-1  ohm,  must  be  used. 


240  COLLEGE    CHEMISTRY 

surface  of  the  acid  very  gently  with  a  thin  glass  rod,  the  ampere- 
meter instantly  responds,  showing  an  increase  in  conductivity.  As 
we  stir,  the  conductivity  increases,  and  the  increase  ceases  only 
when  the  liquid  has  become  homogeneous.  Introduction  of  an 
additional  supply  of  water  will  improve  the  conductivity  still  more, 
but  the  effect  becomes  less  and  less,  until  no  change  on  further 
dilution  is  perceptible.  Reasoning  about  these  effects,  we  perceive 
that  the  amount  of  hydrochloric  acid  has  not  altered  during  the 
experiment.  Yet  the  quantity  of  conducting  material  between  the 
electrodes  must  have  become  greater,  for  the  carrying  power  of 
the  whole  has  improved.  We  were  therefore  observing  the  progress 
of  a  chemical  change  of  the  nonconducting  hydrogen  chloride  into 
conducting  materials.  Hydrogen  chloride  molecules  do  not  carry 
electricity  (p.  145),  but  the  hydrogen  and  the  chloride  ions,  into 
which  it  was  gradually  altered  by  chemical  change  during  the 
stirring,  do  carry  electricity.  Furthermore,  the  change  practically 
ceased  at  great  dilution,  for  the  dissociation  into  ions  was  then 
practically  complete.  If  we  could  conveniently  have  started  with 
only  liquefied,  dry  hydrogen  chloride  in  the  cell,  we  should  have 
observed  the  whole  range  of  changes  from  zero  to  the  maximum. 

When  a  saturated  solution  of  cupric  chloride  is  used  instead  of 
hydrochloric  acid,  dilution  is  accompanied  by  a  similar  improve- 
ment in  conductivity.  Here  we  notice,  besides,  that  the  yellowish- 
green  liquid,  with  which  we  start,  changes  to  a  pale  blue,  as  the 
yellowish-brown  molecules  of  cupric  chloride  are  dissociated  and 
the  color  of  the  solution  becomes  more  exclusively  that  of  the 
copper  ions.  When  the  solution  has  become  perfectly  blue,  further 
dilution  is  seen  to  affect  the  conductivity  but  slightly. 

Reasoning  still  further  about  these  phenomena  we  see  that,  if  we 
start  with  a  fixed  amount  of  a  given  substance,  the  conductivities 
at  different  stages  of  the  dilution  must  be  proportional  to  the  numbers 
of  ions,  and  the  maximum  conductivity  attainable  by  great  dilu- 
tion must  represent  the  effect  when  the  whole  material  has  become 
ionic.  Thus,  if  the  conductivity  at  the  maximum  is  represented, 
say,  by  5,  then  at  the  dilution  where  the  conductivity  is  2,  the 
proportion  of  the  whole  which  is  ionized  is  2/5.  Wlien  the  con- 
ductivity becomes  4,  4/5  of  the  molecules  are  dissociated  and  the 
degree  of  ionization  is  0.8.  When  the  conductivity  becomes  5, 
5/5,  or  all,  of  the  molecules  are  dissociated.  For  example,  in 


IONIZATION 


241 


hydrochloric  acid,  if  we  take  the  normal  solution  (p.  124)  containing 
36.5  g.  of  acid  per  liter  as  the  unit  of  concentration,  the  fractions 
ionized  at  various  concentrations  are  as  follows:  ION,  0.17;  N, 
0.78;  N/10,  0.91;  tf/100,  0.96.  Thus,  measurements  of  con- 
ductivity enable  us  to  study  the  ionic  decomposition  of  all  ionogens, 
and  to  state  accurately  the  fraction  ionized,  at  each  concentration, 
in  solutions  of  every  ionogen.  This  information  is  obviously  most 
valuable,  for  it  places  us  in  a  position  to  know  the  exact  constitu- 
tion of  every  solution  we  use  in  the  laboratory.  In  the  following 
section  the  data  on  which  such  knowledge  can  be  based  is  given. 
In  the  next  chapter  the  mode  of  applying  the  data  is  explained. 

Constitution  of  Solutions  of  lonogens:  Fractions  Ionized. 

—  The  dilute  acids  used  in  the  laboratory  are  generally  of  six  times 
normal  (QN)  concentration.  But,  often,  we  add  only  a  drop  or 
two  to  a  large  bulk  of  liquid,  so  that  the  acids  are  commonly  very 
dilute  as  actually  employed.  The  solutions  of  salts  are  of  different 
strengths,  but  the  great  majority  are  of  normal  (N),  or  even  smaller 
concentrations.  In  practice  they,  also,  are  still  further  consider- 
ably diluted  before  use.  If,  therefore,  we  give  the  fractions  ionized 
(total  molecules  of  ionogen  =  1)  in  decinormal  solutions  (except 
where  otherwise  specified),  the  reader  will  be  able  to  estimate 
roughly  the  proportion  of  each  kind  of  ions  in  any  application  of 
the  reagent.  In  the  case  of  acids  containing  more  than  one  dis- 
placeable  hydrogen  unit,  the  kind  of  ionization  on  which  the  figure 
is  based  is  indicated  by  a  period.  Thus  H.HCOs  means  that  the 
whole  of  the  ionization  is  assumed  to  be  into  H+  and  HCOa". 


FRACTION  IONIZED 


IN  0.1AT  SOLUTIONS  AT  18° 
Acros 


Nitric  acid 0.92 

Nitric  acid  (cone.,  62%)    .    .0.09 

Hydrochloric  acid 0 . 92 

Hydrochloric     acid      (cone., 

35% 0.13 

Sulphuric  acid,  H.H.SO4    .    .  0.61 
Sulphuric  acid  (cone.,  95%).   0.01 

Hydrofluoric  acid 0.15 

Oxalic  acid,  H.HC2O4     .    .    -  0.50 
Tartaric  acid,  H.HT  .    .    .    .0.08 

Acetic  acid  (N) 0.004 

Acetic  acid  .  ....  0.013 


Carbonic  acid,  H.HCO3  . 
Carbonic  acid  (AT/25)  .  . 
Hydrogen  sulphide,  H.HS 
Boric  acid,  H.H2BO3  .  .  . 
Hydrocyanic  acid  .  .  .  . 
Permanganic  acid  (N/2)  . 
Hydriodic  acid  (AT/2)  .  . 
Hydrobromic  acid  (N/2)  . 
Perchloric  acid  (N/2)  .  . 
Chloric  acid  (N/2)  .  .  . 
Phosphoric  acid,  H.H2PO4 
Water  .  .  .  . 


0.0017 

0.0021 

0.0007 

0.0001 

0.0001 

0.93 

0.90 

0.90 

0.88 

0.88 

0.27 

O.Oel 


242 


COLLEGE    CHEMISTRY 


BASES 


Potassium  hydroxide      .    .    .  0.91 

Sodium  hydroxide 0.91 

Barium  hydroxide 0 . 77 

Lithium  hydroxide  (AT)      .    .  0.63 
Tetramethylammonium    hy- 
droxide (N/ 16) 0.96 


Ammonium  hydroxide  .  .  . 
Strontium  hydroxide  (AT/64) 
Barium  hydroxide  (AT/64)  . 
Calcium  hydroxide  (AT/64)  . 
Silver  hydroxide  (AT/1783)  . 
Water  . 


0  .  013 

0.93 

0.92 

0.90 

0.39 

O.OJ 


SALTS 


Sodium  bicarbonate, 

Na.HCO3 0.78 

Sodium  phosphate,  Na2.HP04  0 . 73 

Sodium  tartrate 0. 

Barium  chloride 0, 

Calcium  sulphate  (AT/100)    .0. 

Cupric  sulphate 0 . 

Silver  nitrate 0 . 

Zinc  sulphate 0.40 

Zinc  chloride 0.73 

Mercuric  chloride    .    .   .     (<0.01) 
Mercuric  cyanide    v  .    .    .    Minute 


. 69 
. 77 
64 
39 
.81 


Potassium  chloride     .    .    .    .0.86 

Potassium  nitrate 0 . 83 

Potassium  acetate 0.83 

Potassium  sulphate    .    .   .    .0.72 
Potassium  carbonate  .   .    .    .(0.71) 

Potassium  chlorate 0.83 

Ammonium  chloride  .    .    .    .  0.85  £ 
Sodium  chloride  (N)  .   .   .    .0.66 
Sodium  chloride  (AT/2)  .   .    .  0.74# 

Sodium  chloride 0.84/ 

Sodium  nitrate 0.83// 

Sodium  acetate 0 . 79 

Sodium  sulphate 0 . 70 


In  addition  to  their  use  in  showing  the  nature  of  the  reagents 
employed  in  the  laboratory  (p.  241),  these  numbers  show  also  to 
-~what  extent  any  pair  of  ionic  substances  will  unite  when  mixed  (see 
pp.  247,  251),  and  they  likewise  indicate  the  chemical  activity  of 
the  ionogens  when  in  solution  (see  next  section). 

Relation  of  lonization  to  Chemical  Activity.  —  These 
tables  may  be  used  for  reference.  The  import  of  the  following 
general  statements,  drawn  from  the  tables,  should  be  memorized: 

1.  Salts,  with  the  exception  of  those  of  mercury,  are  all  well 
ionized.     In  actions  involving  their  ions,  salts  are  therefore  all  of 
the  same  order  of  activity,  for  a  dilute  solution  of  every  salt  contains 
a  large  amount  of  the  ionic  components. 

2.  Acids  show  the  most  extreme  differences  in  their  degrees  of 
ionization.     That  is  to  say  then-  solutions  must  contain  very  differ- 
ent concentrations  of  hydrogen-ion.     Since  their  activity  as  acids 
depends  on  this  substance  (p.  217),  and  the  activity  of  a  substance 
is  proportional  to  its  concentration  (p.  182),  it  follows  that  acids 
will  show  very  great  differences  in  apparent  chemical  activity.     At 


IONIZATION  243 

this  point,  therefore,  we  emerge  from  semi-physical  discussion  of 
the  subject  and  reach  something  of  definite,  practical  application 
in  chemical  work. 

The  data  show  that  acids  may  be  divided  roughly  into  four 
classes  with  different  degrees  of  acidic  activity: 

(a)  The  ionization  in  decinormal  solution  exceeds  70  per  cent; 
e.g.,  nitric  acid  and  hydrochloric  acid.  These  are  the  acids  which 
are  chemically  most  active,  for  their  solutions  contain  a  relatively 
high  concentration  of  hydrogen-ion. 

(6)  The  ionization  is  between  70  and  10  per  cent;  e.g.,  sulphuric 
acid  and  phosphoric  acid.  These  acids  are  noticeably  less  active, 
for  their  solutions  contain  a  lower  concentration  of  hydrogen-ion. 

(c)  The  ionization  is  between  10  and  1  per  cent;  e.g.,  acetic  acid. 
These  are  the  weaker  acids,  for  their  solutions  contain  a  very  small 
concentration  of  hydrogen-ion. 

(d)  The  ionization  is  less  than  1  per  cent;  e.g.,  carbonic  and 
boric  acids.     These  are  the  feeble  acids,  for  their  solutions  contain 
only  a  minute  concentration  of  hydrogen-ion. 

3.   The  bases  show  two  classes: 

(a)  Ionization  high;  e.g.,  potassium  hydroxide.  These  bases 
are  active,  for  their  solutions  contain  a  high  concentration  of 
hydroxide-ion. 

(6)  Ionization  less  than  2  per  cent;  e.g.,  ammonium  hydroxide. 
These  bases  are  weak  on  account  of  the  low  concentration  of 
hydroxide-ion. 

4.~  Water  is  less  ionized  than  any  other  substance  in  the  list.  It 
shows  therefore,  as  we  already  know,  usually  little  or  no  interaction 
with  acids,  bases,  or  salts,  and  hence  is  valuable  as  a  solvent  for 
these  substances.  Its  ions  are  H+  and  OH~,  and  it  is  thus  as  much 
(or  as  little)  an  acid  as  a  base. 

Exercises.  —  1 .  With  solutions  of  the  following  substances, 
state,  (a)  what  will  be  the  products  of  electrolysis,  (6)  whether  each 
is  primary  or  secondary,  and  (c)  how  they  may  be  isolated  in  each 
case:  Potassium  chlorate,  potassium  iodide,  potassium  iodate,  sil- 
ver sulphate,  sodium  peroxide. 

2.  Make  equations  (p.  232)  showing  the  ionic  and  molecular 
materials  in  solutions  of  potassium  bromide,  potassium  bromate, 
sodium  periodate,  aluminium  chloride,  zinc  sulphate.  Mark  the 


244  COLLEGE    CHEMISTRY 

charges  on  the  ions  and  give  the  name  of  each  ionic  substance  (p. 
236). 

3.  Prepare  lists  of  other  anions  and  cations  which  have  been 
encountered,  giving  the  formula  and  number  of  charges  of  elec- 
tricity in  each  case. 

4.  If  the  conductivity  of  sodium  chloride  solution  at  the  maxi- 
mum is  110,  and  at  greater  concentrations  is  as  follows:  N,  74.7; 
AT/10,  92.5;   TV/100,  103,  calculate  the  fraction  ionized  at  each 
concentration. 

5.  If  the  conductivity  of  acetic  acid  solution  at  the  maximum  is 
352,  and  at  greater  concentrations  is  as  follows:    10TV,  0.05;   N, 
1.32;  TV/10, 4.6;  TV/100, 14.3,  calculate  the  fraction  ionized  at  each 
concentration. 

6.  If  1  c.c.  of  dilute  hydrochloric  acid  (6TV)  is  added  to  30  c.c. 
of  an  aqueous  solution,  what  is  the  reacting  concentration  of  the 
acid? 

7.  Classify  all  the  acids  in  the  table  (p.  241)  according  to  the 
four  classes  (p.  243). 


CHAPTER  XIX 
IONIC   SUBSTANCES  AND  THEIR  INTERACTIONS 

IN  this  chapter,  after  enumerating  the  various  classes  of  ionqgens, 
and  the  various  kinds  of  ionic  substances,  we  discuss  the  interactions 
of  the  latter.  We  consider  first  the  relations  of  the  ionic  and  the 
molecular  substances  (in  equilibrium)  when  a  single  ionogen  is 
present,  and  then  take  up  the  ways  in  which  such  an  ionic  equi- 
librium is  displaced.  Finally,  we  discuss  some  of  the  useful  ionic 
interactions,  in  which  the  equilibria  are  displaced  so  far  that 
practically  complete  interaction  occurs:  namely,  precipitation, 
neutralization,  and  displacement. 

The  Classes  of  lonogens.  —  Acids  are  classified  according 
to  the  number  of  hydrogen  units  in  their  molecules.  Thus  chloric 
acid  HClOs  is  a  monobasic  acid,  sulphuric  acid  H2S04  a  dibasic  acid, 
and  phosphoric  acid  H3PO4  a  tribasic  acid.  These  terms  relate  to 
the  fact  that,  in  neutralization  (see  p.  254)  the  acids  interact  with 
one,  two,  or  three  molecules  of  a  base  like  sodium  hydroxide. 

Bases  are  named  in  a  similar  way:  sodium  hydroxide  NaOH  is  a 
monoacid  base,  calcium  hydroxide  Ca(OH)2  is  a  diacid  base. 

Salts  like  KC1  and  N^COs  are  neutral  (see  acid  salts,  below)  or 
normal  salts,  and  NaKC03  and  Ca(OCl)Cl  (bleaching  powder)  are 
mixed  salts. 

The  most  interesting  classes  of  mixed  salts  are  the  acid  salts 
(p.  206)  and  the  basic  salts.  In  acid  salts,  like  NaHS04  (p.  141)  and 
KH2PO4  (p.  196),  all  the  hydrogen  of  the  acid  has  not  been  replaced 
by  a  metal.  In  basic  salts,  like  Ca(OH)Cl,  part  of  the  basic 
hydroxyl  remains. 

There  are  also  many  double  salts,  like  ferrous-ammonium  sul- 
phate (NH4)2SO4,FeS04,6H20,  and  alum  (see  index),  some  of  which 
are  in  common  use. 

All  these  substances  are  ionogens  (p.  236).  The  mixed  and 
double  salts  are,  naturally,  dissociated  into  more  than  two  ionic 

substances. 

245 


246  COLLEGE   CHEMISTRY 

Ionic  Substances  Furnished  by  Acids.  —  The  mode  of 
naming  ionic  substances  has  already  been  given  (p.  236). 

Acids,  e.g.,  HC1,  H2S04,  when  dissolved  in  water,  all  furnish 
hydrogen-ion  H*  and  a  negative  ionic  substance  (anion),  e.g.,  Cl~, 
S04=.  The  solutions  differ  from  those  of  salts  in  the  constant  pres- 
ence of  hydrogen-ion,  and  in  the  absence  of  any  other  positive  ion. 

Hydrogen-ion  H+  is  a  colorless  substance.  It  is  sour  in  taste, 
and  its  presence  is  recognized  by  the  fact  that  it  turns  blue  litmus 
red  (see  Indicators,  below).  These  properties  serve  as  tests  for 
acids,  as  they  are  not  interfered  with  by  other  ionic  substances 
which  may  be  present.  Hydrogen-ion  is  univalent  and,  when 
combined  with  negative  radicals  of  salts,  gives  the  (molecular) 
acids.  The  activity  of  acids  depends  upon  the  concentration  of 
the  hydrogen-ion  they  furnish  (p.  242),  and  therefore  upon  their 
solubility  and  the  degree  of  ionization  of  the  dissolved  molecules. 
Some  furnish  so  little  hydrogen-ion  that  their  action  on  litmus 
can  hardly  be  detected. 

Ionic  Substances  Furnished  by  Bases.  —  Bases,  e.g.,  KOH, 
NHiOH,  Zn(OH)2,  all  furnish  hydroxide-ion  OH~  and  some  positive 
ionic  substance  (cation),  K+,  NH4+,  Zn4^.  Their  solutions  differ 
from  those  of  salts  in  the  constant  presence  of  hydroxide-ion  and 
in  the  absence  of  any  other  anion.  The  more  active  bases,  that  is, 
those  which  are  soluble  and  highly  dissociated,  so  that  they  give  a 
high  concentration  of  hydroxide-ion,  are  called  alkalies.  Such  are 
potassium  and  sodium  hydroxides.  They  are  often  named  caustic 
alkalies  and,  individually,  caustic  potash  and  caustic  soda.  The 
solutions  are  called  lyes. 

Hydroxide-ion  OH~  is  a  colorless  substance.  Properties  which 
serve  as  tests  for  bases  are  that  hydroxide-ion  possesses  a  soapy 
taste  and  feeling  and  turns  red  litmus  blue  (see  Indicators,  below) . 
It  is  univalent,  and  combines  with  positive  radicals  to  form 
(molecular)  bases. 

Ionic  Substances  Furnished  by  Salts.  —  Salts  furnish 
positive  and  negative  ionic  substances,  which  may  be  either  simple 
or  composite,  Na.Cl,  Na.NO3,  NH4.C1,  NI^.NOs.  Some  ionic 
substances  are  colored,  Cu**  (cupric-ion)  blue,  CT+++  reddish- 
violet,  Co++  pink,  Mn04~  (permanganate-ion)  purple,  Cr2O7=  (di- 


IONIC    SUBSTANCES   AND   THEIR   INTERACTIONS          247 

chromate-ion)  orange,  but  most  of  them  are  colorless,  K+,  Na+, 
Zn++,  Cl~,  I~,  N03~,  S04~.  They  vary  in  taste,  some  being  salt, 
some  astringent,  some  bitter.  The  ionic  materials  characteristic 
of  salts  do  not  affect  litmus,  and  individual  tests  are  required  for 
each.  Usually  we  add  some  other  ionic  substance,  with  which  the 
ion  thought  to  be  present  combines  to  form  an  insoluble,  molecular 
substance  of  known  color,  or  appearance,  and  examine  the  precipi- 
tate if  any  appears.  Thus,  when  the  presence  of  chloride-ion  Cl~ 
is  suspected,  we  may  add  a  solution  containing  silver-ion  Ag+, 
expecting  to  obtain  a  precipitate  of  silver  chloride  AgCl  (Cl~  + 
Ag+— >  AgClJ).  In  dilute  solutions  of  salts,  the  ions  are  almost 
always  numerous  in  comparison  with  the  molecules  (p.  242),  so 
that  salts  are  practically  all  active  and  their  solutions  almost 
always  respond  readily  to  the  tests  for  the  ions  they  contain. 
The  art  of  detecting  the  various  ionic  substances  present  in  a 
solution  constitutes  a  large  part  of  the  branch  of  chemistry  called 
qualitative  analysis. 

All  the  known  ionic  substances  are  found  in  solutions  of  salts. 
The  only  ions  which  are  not  characteristic  of  salts,  although  some- 
times occurring  in  their  solutions  (see  acid  and  basic  salts,  above) , 
are  hydrogen-ion  H+,  and  hydroxide-ion  OH—. 

It  will  assist  the  reader  if  the  following  facts  are  kept  in  mind. 
The  elements  which  can  form  a  simple  positive  ion  are  the 
metallic  elements  (p.  94,  and  see  Chaps.  XXII  and  XXXIII). 
Non-metallic  elements,  like  nitrogen,  may  be  present  in  a  positive 
ion,  as  in  NH4+,  but  never  exclusively.  In  other  words,  we  know 
no  such  substances  as  nitrogen  sulphate,  or  carbon  nitrate.  Con- 
versely, the  metals  are  frequently  found  in  the  negative  ion,  but 
never  constitute  it  exclusively.  They  are  then  usually  associated 
with  oxygen,  as  in  Mn04~,  and  Cr2O7=. 

The  Ionic  Equilibrium  with  a  Single  lonogen.  —  In  the 

ionization  of  a  molecular  substance,  the  chemical  change  is  incom- 
plete and  the  system  reaches  a  condition  of  equilibrium  (p.  238). 
The  action  is,  therefore,  reversible,  and  there  are  thus  two  routes  to 
the  same  equilibrium  point.  This  fact  must  not  be  forgotten,  for 
we  have  to  consider  the  union  of  ionic  substances  even  more  often 
than  the  converse  change.  Now,  the  degrees  of  ionization  of  various 
ionogens  tell  us  the  location  of  the  equilibrium  point,  and  therefore 


248  COLLEGE    CHEMISTRY 

the  extent  of  the  chemical  change  involved  in  reaching  this  point  by 
either  route,  that  is,  either  by  the  dissociation  of  molecules  or  by  the 
union  of  ions.  In  a  class  of  interactions,  of  which  all  are  incom- 
plete, and  only  those  are  interesting  and  useful  which  approach 
completeness,  we  require  some  means  of  knowing  which  are  com- 
plete and  why  they  are  so.  The  table  of  fractions  ionized  (p.  241) 
supplies  most  of  the  required  information. 

To  illustrate,  take  the  case  of  a  single  ionogen.  When  we  place 
hydrogen  chloride  in  decinormal  solution,  0.92  of  the  molecules  dis- 
sociate. Conversely,  when  we  start  with  the  hydrogen-ion  and 
chloride-ion,  say  by  mixing  two  solutions,  each  of  which  contains 
one  of  them  (along  with  another  ion),  then  1  —  0.92,  or  only  0.08 
of  these  ionic  substances  will  combine. 

This  exemplifies  the  case  of  an  active  acid.  The  following  equa- 
tions show  the  data  for  six  typical  substances  in  N/10  solution, 
namely,  two  acids,  two  bases,  and  two  salts: 

(8%)HC1    ^H+  +Cl-(92%),    (98.7%)HC2HA^H+     +C2H30!i-(1.3%) 
(9%)KOH  <=±  K+  +  OH-(91  %),   (98.7%)NH4OH  +±  NIL+  +  OH-(1.3%) 
(16%)NaCl  <±  Na+  +  Cl-(84%),          (61  %)CuSO4  ^  CirH-  +  80^(39%) 

These  samples  are  chosen  to  illustrate,  in  each  pair,  the  extremes. 
Thus,  when  potassium-ion  and  hydroxide-ion  are  brought  together 
little  union  takes  place,  while  with  ammonium-ion  and  hydroxide- 
ion  the  union  is  practically  complete.  In  the  case  of  the  soluble 
salts,  however,  there  are  almost  (p.  242)  no  cases  of  considerable 
union  of  the  ions  in  dilute  solutions.  The  case  of  water,  on  the 
other  hand,  is  one  of  the  most  extreme: 

(99.96%)  H20  <=±  H+  +  OH' 


Hydroxide-ion  and  hydrogen-ion  thus  unite  almost  completely. 

Similar  reasoning  enables  us  to  handle  the  more  complex,  but 
very  common  case  of  the  mixing  of  two  ionogens.  The  degrees  of 
ionization  tell  us  the  exact  condition  of  each  system  separately, 
before  mixing.  The  result  of  the  mixing  is  best  understood  by 
viewing  the  change  as  consisting  in  a  displacement  of  each  of  the 
equilibria  by  the  action  of  the  components  of  the  other.  We  con- 
sider, therefore,  next,  the  displacement  of  ionic  equilibria. 

The  Displacement  of  Ionic  Equilibria.  —  Equilibria  are 
displaced  by  changes  which  favor  or  disfavor  one  of  the  opposed 


IONIC   SUBSTANCES   AND   THEIR   INTERACTIONS          249 

actions  (p.  180).  There  may  be  either,  (1)  a  physical  change  in  the 
conditions,  or  a  chemical  interaction  which  (2)  adds  to,  or  (3) 
removes  one  of  the  interacting  substances.  Each  of  these  may  be 
illustrated  in  turn. 

1.  As  an  example  of  the  first,  we  have  the  effect  of  changing  the 
amount  of  the  solvent  (p.  215) .  Adding  more  of  the  solvent  reduces 
the  concentration  of  the  ionic  materials  and  disfavors  their  union, 
so  that  it  indirectly  promotes  dissociation.  The  larger  the  volume 
in  which  the  ions  are  scattered,  the  less  often  will  they  meet,  and 
the  smaller  the  amount  of  combination.  On  the  other  hand, 
evaporating  off  a  part  of  the  solvent  favors  the  encounters  of  the 
ions  and  promotes  combination.  When  the  solvent  is  at  last 
entirely  gone,  the  whole  material  is  molecular. 

In  cases  where  the  ionic  and  molecular  substances  are  all  color- 
less, these  changes  can  be  followed  only  by  a  study  of  the  freezing- 
points  or  other  similar  properties  of  the  solutions  (p.  216).  But 
when  the  substances  are  of  different  colors,  the  changes  can  also  be 
seen.  Thus,  cupric  bromide  in  the  solid  form  is  a  jet  black,  shining, 
crystalline  substance.  When  treated  with  a  small  amount  of 
water  it  forms  a  solution  which  is  of  a  deep  reddish-brown  tint, 
giving  no  hint  of  resemblance  to  a  solution  of  any  cupric  salt.  This 
doubtless  represents  the  color  of  the  molecules.  When  more  water 
is  added,  the  deep  brown  gives  place  gradually  to  green,  and  finally 
to  blue.  The  latter  is  the  color  of  the  cupric-ion  (Cu++),  and  is 
familiar  in  all  solutions  of  cupric  salts.  The  colorless  nature  of 
solutions  of  potassium  and  sodium  bromides  shows  that  bromide- 
ion  (Br~)  is  without  color.  Hence,  in  the  present  instance  it  is 
invisible.  We  are  thus  watching  the  forward  displacement  of  the 
equilibrium: 

CuBr2  (brown)  *=;  GU++  (blue)  +  2Br~. 

If  1  g.  of  the  solid  is  taken,  it  dissolves  in  about  its  own  weight  of 
water,  and  independent  measurement  shows  that  there  is  relatively 
little  ionization.  Hence  the  solution  is  deep  brown.  When  10  c.c. 
of  water  has  been  added,  70  per  cent  of  the  salt  is  ionized,  and  the 
solution  is  green.  With  40  c.c.  of  water,  only  19  per  cent  remains 
in  molecular  form,  and  the  blue  color  of  the  cupric-ion  entirely 
overbears  the  tint  of  the  molecules.  If  we  now  remove  the  water 
by  evaporation,  all  these  changes  are  reversed.  When  30  c.c.  of 


250  COLLEGE    CHEMISTRY 

the  water  has  been  driven  off,  the  solution  is  green.  As  the  evapo- 
ration-of  the  remaining  10  c.c.  progresses,  the  brown  color  appears. 
When  the  water  is  all  gone,  the  black  residue  remains.  Here  we 
are  observing  the  backward  displacement  of  the  equilibrium, 
CuBr2  ±5  Cu++  +  2Br~. 

2.  Cupric  bromide  may  be  used  to  illustrate  also  the  chemical 
methods  of  displacing  equilibria.     Thus,  we  may  show  the  effect  of 
adding  more  of  one  of  the  reacting  substances.     If,  at  the  green  stage, 
we  dissolve  solid  potassium  bromide  in  the  liquid  (KBr<=±K++Br~), 
the  increased  concentration  of  bromide-ion  causes  more  vigorous 
interaction  of  the  ions,  and  the  molecules,  with  their  brown  color, 
become  prominent  again.     Adding  cupric  chloride  increases  the 
concentration  of  cupric-ion  and  has  the  same  effect.     In  either 
case,  renewed  dilution  with  water  reduces  the  concentrations  of  all 
the  ions  once  more,  the  molecules  become  fewer,  and  the  brown 
color  is  displaced  by  the  blue  for  the  second  time. 

3.  Finally,  the  displacement  of  the  same  equilibrium  by  remov- 
ing one  of  the  interacting  substances  may  be  illustrated.     Thus,  if 
the  chocolate-brown  solution,  in  which  molecular  cupric  bromide 
predominates,  is  shaken  with  pulverized  lead  nitrate  (and  filtered), 
two  changes  are  noticed.     A  pale  yellow  precipitate  of  lead  bromide 
appears  (Pb4"1"  +  2Br~  — » PbBr2  J, ),  and  the  brown  color  fades  into 
green.     Here  the  displacement  is  the  opposite  of  the  last.     Instead 
of  reinforcing  one  of  the  ions,  we  have  reduced  the  concentration, 
and  in  fact  almost  entirely  removed  one  of  them,  namely  Br~. 
This  has,  naturally,  stopped  the  interaction  of  the  Cu++  and  Br~ 
which  reproduces  the  brown,  molecular  CuBr2.     Hence  the  disso- 
ciation of  the  latter  has  continued  to  exhaustion  of  the  whole 
molecular  material. 

The  reader  will  find  that  the  behavior  of  these  ionic  equilibria, 
and  the  way  in  which  we  discuss  and  explain  it,  are  complete 
parallels  of  the  behavior  and  explanation  in  the  case  of  ordinary 
equilibria  (pp.  185-187),  which  should  now  be  reexamined.  The 
illustrations  in  the  present  section,  and  particularly  the  third  (cf. 
p.  203),  should  be  considered  until  every  feature  is  perfectly  clear. 
They  furnish  the  key  to  understanding  the  applications  which  fol- 
low. One  fact  must  not  escape  notice,  and  that  is  that  in  none 
of  the  three  instances  was  the  forward  action  (the  dissociation) 
in  itself  affected.  The  molecules  of  cupric  bromide  have,  as  we 


IONIC   SUBSTANCES  AND   THEIB   INTERACTIONS         251 

should  expect,  a  certain  tendency  to  decompose.  No  encounters 
between  these  molecules  are  required  for  mere  decomposition. 
Hence  their  decomposition  is  not  influenced  by  their  nearness  to, 
or  remoteness  from,  one  another  (illustration  1),  nor  by  the  presence 
of  any  other  kinds  of  molecules  or  ions  (illustrations  2  and  3). 
The  effect,  whether  it  involved  an  apparent  increase,  or  a  diminu- 
tion of  the  dissociation,  was  always  accomplished  by  altering  the 
concentration  of  the  ionic  substances,  and  therefore  the  activity  of  the 
reverse  action. 

Applications:   Double  Decomposition  in  Solution.  —  We 

are  now  prepared  to  consider  the  general  case  of  mixing  the  solu- 
tions of  two  ionogens. 

When  solutions  of  two  ionized  substances  are  mixed,  the  first  reflec- 
tion which  occurs  to  us  is  that  each  of  these  has  been  diluted  by  the 
water  in  which  the  other  was  dissolved,  so  that  the  first  effect  will 
be  to  increase  the  degree  of  ionization  of  both  to  a  certain  extent. 

The  next  consideration  is,  however,  that  we  have  produced  a 
mixture  of  four  ions,  which  must  have  at  least  some  tendency  to 
unite  crosswise.  Thus  potassium  chloride  and  sodium  nitrate  in 
dilute  solution  are  very  greatly  ionized  before  mixing.  The  re- 
versible actions,  represented  by  the  horizontal  pair  of  the  following 
equations,  have  taken  place  extensively.  But,  by  mixing  the 
liquids,  we  have  brought  into  presence  of  -^Q  4__  -^+  ,  Q- 
one  another  two  new  pairs  of  positive  and  NaNO  4—  J^-Q  -_i_  jq-a+ 
negative  ions.  Hence,  two  other  reversi-  *  •* 

ble  actions,  the  vertical  ones,  will  be  set  KNO       NaCl 

up  and  will  proceed  until  a  fresh  equi- 
librium of  all  the  ions  with  all  four  kinds  of  molecules  has  been 
reached.      Thus  far  the  description  will  fit  any  case  of  mixing 
solutions  of  two  ionogens. 

Now,  in  this  particular  instance,  what  is  the  actual  extent  of  such 
interaction  as  has  occurred?  To  answer  this  question  we  require  to 
know  the  proportion  of  molecules  to  ions  in  a  solution  of  each  of  the 
four  salts  (p.  242).  In  decinormal  solutions  it  is  KC1,  14  : 86; 
NaN03,  17  :  83;  KN03,  17  : 83,  NaCl,  16  :  84,  so  that  the  salts 
are  all  equally  well  ionized.  It  is  a  good  plan  to  add  these  pro- 
portions in  the  formulation.  Furthermore,  in  a  dilute  mixture, 
such  as  we  are  considering,  the  proportions  of  ions  are  greater  than 


252  COLLEGE    CHEMISTEY 

these  figures  indicate.     Hence,  practically  no  chemical  action  has 
occurred. 


(14%)KC1  ±=>  K+  +      OP  (86%) 

(17%)NaN03  ±=*  NO3~  +      Na+  (83%) 

JT  It 

KN03  NaCl 


That  this  inference  is  correct  is  shown  by  independent  evidence. 
Thus  when  the  solutions  of  salts  are  mixed,  no  thermal  effect  is 
observable.  This  fact  has  been  known  since  1842  as  Hess'  law 
of  thermoneutrality.  Again,  if  the  solutions  are  placed  in  a  cell 
(Fig.  81,  p.  239),  so  that  the  one  forms  a  layer  below  the  other,  no 
change  in  conductivity  is  noticed  when  the  solutions  are  stirred 
together.  Hence  no  change  in  the  number  of  ions  has  occurred. 

We  conclude,  then,  that  when  two  highly  ionized  substances  are 
mixed,  and  the  possible  products  are  also  highly  ionized,  soluble 
substances,  then  practically  no  chemical  action  occurs.  This  rule 
applies  to  dilute  solutions  of  all  soluble  salts  (p.  242)  and  to  mixing 
salts  with  the  highly  ionized  acids  or  bases. 

Conversely,  when  two  ionized  substances  are  mixed,  an  extensive 
chemical  change  does  ensue  in  two  cases,  namely  : 

1.  When  one  of  the  possible  products  is  an  insoluble  substance 
and  precipitation  occurs,  for  this  removes  the  ions  used  to  form  the 
insoluble  body. 

2.  When  one  of  the  possible  products,  although  soluble,  is  little 
ionized,  as  in  neutralization,  for  this  likewise  removes  the  ions  re- 
quired to  form  molecules  of  the  product.     We  proceed,  therefore, 
to  discuss  these  two  important  classes  of  actions. 

Precipitation.  —  A  typical  case  of  precipitation  occurs  when 
we  mix  dilute  solutions  of  silver  nitrate  and  sodium  chloride. 


(16%)  NaCl  fc?  Na+    +  Cl~  (84%) 
(19%)  AgN03  fc*  NOT,  +  Ag+  (81%) 

Jt  Jf 

NaN03     AgCl  (dslvd) 


AgCl  (solid) 


IONIC    SUBSTANCES   AND   THEIR  INTERACTIONS          253 

Here,  since  the  four  substances  are  all  salts,  they  are  all  highly 
ionized.  If  they  were  all  soluble,  then,  in  dilute  solutions,  perhaps 
5  per  cent  of  each  salt  would  be  in  molecules  and  the  rest  in  ionic 
form.  But  the  molecules  of  silver  chloride  are  excessively  insoluble. 
In  all  cases  of  precipitation,  we  look  up  the  solubilities  of  the  possible 
products  (see  Table  of  Solubilities  inside  the  front  cover).  Here 
we  find  that  one  liter  of  water  will  dissolve  only  0.0016  g.  silver 
chloride  (this  quantity  includes  both  ions  and  molecules).  So 
the  concentration  of  the  AgCl  (dslvd)  becomes  almost  zero  through 
precipitation.  So  far  as  it  is  in  solution,  however,  being  a  salt 
and  very  dilute,  it  is  practically  all  ionized.  The  precipitation 
displaces  the  equilibrium,  for,  the  dissociation  having  thus  ceased, 
those  of  the  ions  Ag+  and  Cl~  which  combine  are  not  replaced  by 
others.  Hence  the  silver-ion  and  chloride-ion  almost  disappear. 
This  occurrence  affects  in  turn  the  equilibria  with  Na+  and  NO3~, 
so  that  the  NaCl  and  AgN03  become  completely  ionized.  Hence 
the  concentrations  of  NaCl  and  AgN03,  of  Ag+  and  Cl~,  and  of 
the  dissolved  AgCl,  all  become  practically  zero  at  last.  The 
system  finally  contains  only  a  precipitate  of  molecular,  solid  silver 
chloride  and  a  solution  of  the  three  substances,  Na+  +  NO3~  *=? 
NaN03,  in  equilibrium.  By  far  the  greater  part  of  this  material 
in  solution  is  the  ionic,  namely  the  Na+  and  the  N03  . 

To  avoid  a  misconception,  note  that  the  answer  to  the  question, 
"Is  silver  chloride  a  highly  ionized  substance?"  is  "Yes."  Since 
it  is  a  salt,  we  expect  this.  True,  very  little  of  it  dissolves,  so  that 
it  cannot  give  many  ions  to  a  solution.  But  little  or  much  ionized 
refers  to  the  proportion  ionized  of  the  material  which  has  dissolved. 
With  undissolved  material  ionization  has  nothing  to  do. 

It  should  be  noted  that,  when  the  solutions  are  mixed,  as  in  the 
foregoing  example,  strictly  speaking,  the  chief  interaction  taking 
place  is  the  production  of  the  insoluble  body.     The  largest  part  of 
the  chemical  action  may  be  formulated  thus: 
Agf  +  cr  -» AgCl. 

The  chief  change  that  has  as  yet  befallen  the  ions  of  sodium  nitrate 
is  that  they  have  been  transferred  from  two  separate  vessels  into 
one.  Potentially  the  salt  has  been  formed.  But  the  actual  union 
of  its  ions,  to  give  the  second  product  in  the  molecular  condition, 
Na+  +  NOr  ->  NaN03, 


254  COLLEGE    CHEMISTRY 

comes  about  only  when,  at  some  subsequent  time,  if  at  all,  the 
water  is  evaporated  away. 

The  foregoing  formulation  and  explanation  apply  to  every  case 
of  mixing  ionogens  where  precipitation  occurs,  that  is,  where  the 
products  are  insoluble  acids,  bases,  or  salts. 

Neutralization.  —  We  may  now  consider  the  case  of  mixing 
solutions  of  two  ionogens  where  one  is  an  acid  and  one  a  base. 

(>87<7)  (  0  1<7)  The  general  Plan  of  a11  in~ 

(8%)  HC1  £,  OP  °+  H+     (92%)  factions  of  aci^  and  bases 

(9%)  NaOH^Na+  +  OH~  (91%)  *  shown  m.the  fo™uktion. 

•I            |*  The  lomzation  of  the  hydro- 

<*<*  m  chloric  acid  reaches  0.92  in  a 

NaCl       H2O  ,    . 

(<?  1  W  ^    n  one/  ^          decmormal  solution,  and  goes 
farther  when  the  acid  is  di- 

luted  with  the  water  of  another  solution.  That  of  the  sodium  hy- 
droxide similarly  goes  beyond  0.91.  Thus  the  substances  in  the 
solutions  before  mixing  are  almost  entirely  ionic.  The  crosswise 
union,  H+  +  OH~  *=>  H2O,  however,  is  all  but  complete,  for  water 
i&  hardly  ionized  at  all  (p.  243).  The  materials  on  whose  inter- 
action with  the  Cl~  and  Na+,  respectively,  the  maintenance  of  the 
molecules  HC1  and  NaOH  depends,  being  thus  removed,  the  disso- 
ciation of  the  acid  and  base  promptly  brings  itself  to  completion, 
and  the  left  sides  of  the  equations  vanish.  Practically  all  the 
hydrogen-ion  and  hydroxide-ion  become  water,  which  thenceforth 
is  simply  a  part  of  the  solvent.  The  Cl~  and  Na+,  however,  if 
the  solution  is  now  1/20  normal,  unite  to  the  extent  of  0.13  only. 
If  it  is  more  dilute,  this  union  forms  a  still  smaller  factor  in  the 
whole  change.  Practically  it  is  negligible.  Now  all  that  has  been 
said  of  this  acid  and  base  will  apply  mutatis  mutandis  whenever 
any  active,  highly  ionized  acid  and  base  come  together.  Thus 
we  may  write  one  simple  equation  for  all  neutralizations  of  active 
acids  and  bases: 


without  omitting  anything  essential. 

The  ions  of  a  salt  are  always  left  over  from  the  main  action,  and 
may  be  brought  together,  in  turn,  by  evaporation:  Na++Cl~—  » 
NaCl,  or  the  liquid  may  be  used  as  a  solution  of  the  pure  salt. 


IONIC    SUBSTANCES   AND   THEIR   INTERACTIONS          255 

Confirmations  of  this  View  of  Neutralization.  —  That 
these  inferences  are  correct  is  shown  by  many  facts.  The  most 
conspicuous  of  these  is  the  fact  that,  when  equivalent  amounts  of 
active  acids  and  bases  are  used,  the  mixture  is  without  action  either 
on  red  or  on  blue  litmus.  It  is  neutral  to  indicators  —  hence  the 
term  neutralization  applied  to  the  operation  of  mixing  an  acid  and  a 
base.  Specifically,  the  absence  of  effect  upon  litmus  demonstrates 
the  absence  of  hydrogen-ion  H+  and  of  hydroxide-ion  OH~,  alike, 
in  the  product,  and  confirms  the  theory. 

Again,  a  considerable  thermal  effect  accompanies  neutralization. 
But,  in  the  cases  we  are  discussing,  that  is  where  active  bases  and 
acids  are  employed,  the  heat  liberated  by  use  of  equivalent  weights 
(p.  124)  is  always  the  same,  namely  13,700  cal.  That  it  is  always 
the  same  confirms  our  theory,  for  practically  the  whole  change  is 
always  the  formation  of  18  g.  of  water  from  the  ions. 

Still  again,  when  we  place  the  acid  and  base  in  the  cell  (Fig.  81, 
p.  239),  so  that  the  one  forms  a  layer  beneath  the  other,  and  watch 
the  amperemeter  while  we  mix  the  solutions,  a  marked  decrease  in 
the  current  passing  through  the  cell  is  noticed.  This  also  confirms 
our  theory,  for  it  is  our  belief  that  one-half  of  the  ions,  namely  the 
H+  and  OH~,  disappear  as  such  during  the  action.  The  decrease 
is,  in  fact,  to  less  than  half  the  reading  before  mixing,  because  the 
two  speediest  ions  have  been  removed. 

When  less  highly  ionized  acids  or  bases  are  used,  the  only  differ- 
ence is  that  there  are  more  of  the  molecular  materials  present, 
before  the  solutions  are  mixed.  But  the  removal  of  the  H+  and 
OH~  ions  permits  the  molecules  of  the  acid  and  base  to  dissociate, 
so  that  the  final  products  are  water  and  the  ions  of  a  salt,  as  before. 

The  foregoing  formulation  and  explanation  apply  to  every  case  of 
mixing  ionogens,  where  a  very  slightly  ionized  substance  is  one  of 
the  products,  that  is,  when  water,  or  a  feeble  acid,  or  a  feeble  base 
(pp.  242-243)  is  formed. 

Acidimetry  and  Alkalimetry.  —  When,  as  is  constantly  the 
case,  a  chemist  desires  to  ascertain  the  quantity  of  an  acid  or  base 
present  in  a  solution,  he  uses  for  the  purpose  the  interaction  just 
discussed.  If,  for  example,  the  problem  is  to  ascertain  the  weight 
of  hydrogen  chloride  in  each  liter  of  a  specimen  of  hydrochloric 
acid,  this  can  be  done  by  neutralizing  a  measured  portion  of  this 


256 


COLLEGE    CHEMISTRY 


acid  with  a  solution  of  an  alkali  of  known  concentration.  The 
volume  of  the  latter  which  is  required  for  the  purpose  is  observed. 
If  the  alkali  is  sodium  hydroxide,  the  action  taking  place  is 

HC1  +  NaOH  ->  H2O  +  NaCl. 

The  volume  of  acid  is  measured  out  into  a  beaker  by  means  of  a 
pipette  (Fig.  82)  of  fixed  capacity,  which  is  filled  by  suction  to  the 

mark  on  the  stem.  Sup- 
pose the  amount  to  be 
25  cc.  The  standard 
alkali  solution  is  placed 
in  a  burette  (Fig.  83), 
which  is  filled  down  to 
^—  the  tip  of  the  nozzle.  A 
few  drops  of  litmus  solu- 
tion are  now  added  to 
the  acid,  and  the  alkali  is 
allowed  to  run  in  slowly. 
After  a  time,  the  hy- 
droxide-ion which  this 
introduces  will  begin  to 
produce  a  blue  color, 
close  to  where  the 
stream  enters  the  liquid. 
This  is  at  first  dissi- 
pated by  stirring,  and 
the  whole  remains  red. 
Finally,  however,  a 
point  is  reached  at 
which  the  entire  solu- 
tion assumes  a  tint  in- 
termediate between 
blue  and  red.  With 
one  drop  less  of  the 
base,  it  is  distinctly  red.  With  one  drop  more,  it  would  become 
distinctly  blue.  Litmus  paper  of  either  shade  dipped  in  this  neu- 
tral solution  remains  unaffected. 

By  the  use  of  a  standard  solution  of  an  acid  in  the  burette,  the 
quantity  of  a  base  may  be  determined  in  the  same  way. 


FIG.  82. 


Fio.  83. 


IONIC   SUBSTANCES  AND   THEIR  INTERACTIONS         257 

Standard  Solutions.  —  The  standard  solutions  used  in  this 
work  are  usually  normal,  and  contain  one  equivalent  weight  of  the 
alkali  or  acid  in  one  liter  of  the  solution.  For  more  delicate  work, 
decinormal  (N/10)  solutions  may  be  employed.  The  concentra- 
tion of  such  a  solution  is  called  its  titer,  and  the  operation  of 
analyzing  another  solution  by  means  of  it,  titration.  The  value  of 
standard  solutions  lies  in  the  fact  that,  when  once  the  solution  has 
been  prepared,  and  the  exact  concentration  adjusted  by  quantita- 
tive experiments,  its  use  does  not  require  any  weighing,  and  the 
measurements  of  volumes  can  be  carried  out  with  great  rapidity. 
The  calculation  of  the  result  is  also  simple.  One  liter  of  normal 
alkali  contains  17  g.  of  available  hydroxyl,  and  one  liter  of  normal 
acid,  1  g.  of  available  hydrogen  (p.  124).  Equal  volumes  of 
normal  solutions  will  therefore  exactly  neutralize  one  another,  18  g. 
of  water  being  formed  by  interaction  of  a  liter  of  each.  If,  for 
the  neutralization  of  the  25  c.c.  of  hydrochloric  acid  used  above, 
50  c.c.  of  normal  alkali  are  required,  the  acid  is  twice-normal  (2N). 
When  15  c.c.  are  required,  the  acid  is  if  or  f  JV.  If  the  actual 
weight  of  the  acid  in  the  latter  case  has  to  be  calculated,  we  remem- 
ber that  there  are  36.46  g.  of  hydrogen  chloride  in  1  1.  of  a  normal 
solution,  and  therefore  36.46  X  f  X  rffo  g.  =  0.5467  g.  in  25  c.c. 
of  a  solution  which  is  f-normal. 

Methods  of  quantitative  analysis  in  which  standard  solutions  are 
employed  are  known  as  volumetric  methods,  and  are  much  used  by 
analysts  and  investigators.  They  occupy  much  less  time  than 
gravimetric  operations,  in  which  numerous  weighings  have  to 
made  and  are  often  just  as  accurate.  The  substances  like  litmus, 
S.  whose  change  of  color  the  completeness  of  the  action  is  made 
known,  are  called  indicators. 

Indicators.  -  Indicators  are  substances  which  in  presence  of 
certain  other  substances,  assume  a  very  deep  color or  change 
sharply  from  one  deep  color  to  another.  Thus  phenolphthalem 
is  colorless  in  presence  of  acids  (t.«  hydrogen-ion),  and  red  when 


258  COLLEGE   CHEMISTRY 

indicator  is  so  small  as  to  be  negligible.  The  common  indicators 
are: 

Phenolphthalein  Ci4Hio04,  a  colorless  substance  and  very  feeble 
acid.  It  is  not  perceptibly  dissociated  into  its  ions, 

CuHio04  (colorless)  *=5  Ci4H904~  (red)  +  H+, 

and  in  neutral  or  acid  solutions  is,  therefore,  without  visible  color. 
When  a  base  is  added  gradually  to  an  acid  containing  some  of  this 
indicator,  the  acid  is  first  neutralized.  Then,  and  not  till  then,  the 
slightest  excess  of  hydroxide-ion  unites  with  the  trace  of  hydro- 
gen-ion from  the  phenolphthalem,  the  above  equilibrium  is  dis- 
placed forwards,  and  a  visible  amount  of  the  red  negative  ion 
is  formed: 

Ci4Hio04  (colorless)  ±=?  Ci4H904~  (red)  +  H+      )  <_  H  o 
NaOH  fc?  Na+  +  OH~  J  " 

In  this  more  compact  formulation,  we  show  the  product  (H20) 
from  the  union  of  the  two  ions  which  combine,  but  omit  the  prod- 
uct from  the  union  of  Na+  and  Ci4H9O4~,  because  here  (since  the 
product  is  a  salt)  hardly  any  union  occurs. 

Litmus  is  an  extract  from  certain  lichens,  first  used  by  Boyle. 
It  contains  azolitmin.  One  of  its  colors  is  that  of  the  molecule, 
and  the  other  that  of  the  ion. 

Methyl  orange  (CH3)2NC6H4.N  :  N.Cel^SOsNa  is  a  complex  or- 
ganic compound  which  gives,  in  acid  solution,  a  red,  and  in  alka- 
line solution  a  yellow  color. 

Congo  red  is  the  sodium  salt  of  an  acid  of  complex  structure  (see 
Dyes).  In  neutral  or  alkaline  solutions  it  is  red;  with  acids  it 
turns  blue.  Paper  dipped  in  Congo  red  differs  from  litmus  paper 
in  that  it  shows  gradations  in  color,  the  blue  being  much  more 
distinct  with  an  active  acid  than  with  a  relatively  weak  one  like 
acetic  acid  (p.  241).  Litmus  paper  is  equally  red  with  all  acids 
save  the  very  feeblest. 

Displacement:  The  Electromotive  Series.  —  In  the  preced- 
ing sections  we  have  dealt  with  cases  in  which  ionic  substances 
underwent  combination  or  ionogens  dissociated.  This  is  one  of  five 
kinds  of  ionic  chemical  change.  Of  the  remaining  four,  ionic  dis- 


IONIC   SUBSTANCES   AND   THEIR   INTERACTIONS         259 

placement  is  the  one  *  that  we  have  most  frequently  encountered. 
Thus,  certain  metals  displace  hydrogen  from  dilute  acids  (p.  60)  : 

Zn  +  H2S04  ->  ZnS04  +  H2. 

These  interactions  do  not  occur  in  the  absence  of  water  (p.  53),  and 
now  appear  in  a  new  light,  namely,  as  ionic  actions: 

Zn  +  2H+  +  S04=  ->  Zn++  +  H2  +  S04=. 

The  molecular  sulphuric  acid  and  zinc  sulphate,  which  are  small 
in  amount,  are  omitted  because  they  do  not,  as  such,  take  part  in 
the  change.  On  looking  at  the  equation,  we  perceive  that  the 
sulphate-ion  is  also  unaltered  by  the  action,  and  may  be  left  out 
likewise  : 

Zn  +  2H+  ->  Zn++  +  H2. 

True,  hydrogen-ion  cannot  be  used  alone,  for  it  is  always  accom- 
panied by  some  negative  radical.  But  the  latter,  like  the  vessel  in 
which  the  experiment  is  made,  is  part  of  the  necessary  apparatus, 
and  not  an  interacting  substance.  The  change  has  consisted  in  the 
ionization  of  the  zinc,  and  the  transfer  to  it  of  the  electric  charge 
of  the  hydrogen-ion.  In  terms  of  electrons  (p.  235),  each  atom 
of  zinc  has  lost  two  electrons  (Zn  —  2e  =  Zn++)  and  two  ions  of 
hydrogen  have  taken  up  the  electrons  (2H+  +  2e  —  >  H2). 

These  statements  enable  us  to  understand  why  active  acids,  with 
zinc,  give  hydrogen  faster  than  do  inactive  acids  (p.  54).  The 
former  provide  a  higher  concentration  (p.  243)  of  hydrogen-ion, 
that  is,  of  the  real  interacting  substance,  than  do  the  latter. 

A  similar  displacement  of  negative  ions  has  been  met  with  (pp. 
194,  199).  Thus,  chlorine  displaces  bromine  from  solutions  con- 
taining bromide-ion. 


The  Electromotive  Series.  —  Displacement  occurs  with  all 
positive  ions.  Thus,  zinc  will  displace  other  metallic  elements, 
such  as  iron,  lead,  copper,  and  silver,  from  the  ionic  conditions, 
when  it  is  placed  in  solutions  of  their  salts: 

Zn  +  Cu++  ->  Zn++  -f-  Cu. 

*  The  discharge  of  an  ion  and  liberation  of  its  material  in  electrolysis 
(pp.  55,  155,  227)  is  another.  Attention  will  be  called  to  the  remaining  two 
when  suitable  illustrations  occur  (see  pp.  270,  504). 


260 


COLLEGE   CHEMISTRY 


Here  the  copper  appears  as  a  red  precipitate.  Lead,  in  turn,  will 
displace  copper  and  silver,  but  not  zinc  or  iron.  Copper  will  dis- 
place silver.  Thus  the  metals  can  be  set  down  in  an  order,  such 
that  each  metal  displaces  those  following  it  in  the  list  and  is 
displaced  by  those  preceding  it.  This  list  is  known  as  the  electro- 
motive series  of  the  metals,  because  in  electrolysis  of  normal  solu- 
tions of  their  salts,  the  electromotive  force  of  the 
current  required  to  deposit  each  metal  is  less 
than  that  for  the  metal  preceding  in  the  list. 
For  present  purposes,  the  list  shows  the  metals 
in  the  order  of  diminishing  tendency  to  enter  the 
ionic  from  the  elementary  condition. 

The  electromotive  series  embodies  many  facts 
in  the  behavior  of  the  metals,  and  should  be  kept 
in  mind  as  furnishing  a  key  to  all  actions  in- 
volving solutions  in  which  a  free  metal  is  used 
or  produced.  It  is,  in  fact,  identical  with  the 
order  of  activity  (p.  60). 

To  avoid  a  common  misconception,  it  must 
be  noted  that  the  electromotive  series  cannot  be 
used  to  explain  the  tendency  of  one  radical  to 
dislodge  another  in  double  decompositions.  The 
place  of  an  element  in  the  E.M.  series  defines 
its  relative  activity  when  free,  and  has  to  do  only 
with  actions  where  one  free  element  displaces 
(p.  55)  another.  The  influences  which  deter- 
mine a  double  decomposition  (cf.  pp.  143,  186) 
are  such  as  the  insolubility  of  a  compound. 
Thus,  potassium  bromide  solution  will  slowly 
convert  a  precipitate  of  silver  chloride  into  one 
of  silver  bromide:  AgCl  +  KBr  -»  AgBr  -f-  KC1. 
This  occurs  because  silver  bromide  is  the  less  soluble  salt.  But/ree 
bromine  never  displaces  chlorine  from  binary  combination  with  a 
metallic  element.  It  is  free  chlorine  that  displaces  combined 
bromine. 

Non-Ionic  Modes  of  Forming  lonogens.  —  While  ionogens 
may  always  be  made  by  the  union  of  the  proper  ions,  they  must 
nevertheless,  in  the  absence  of  the  solvent,  be  regarded  as  chemical 


ELECTROMOTIVE 

SERIES  OF  THE 

METALS. 

Potassium 

Sodium 

Barium 

Strontium 

Calcium 

Magnesium 

Aluminium 

Manganese 

Zinc 

Chromium 

Cadmium 

Iron 

Cobalt 

Nickel 

Tin 

Lead 

Hydrogen 

Copper 

Arsenic 

Bismuth 

Antimony 

Mercury 

Silver 

Palladium 

Platinum 

Gold 


IONIC    SUBSTANCES   AND   THEIR   INTERACTIONS          261 

substances  which  may  be  constructed,  and  very  frequently  are 
made,  out  of  their  constituents  without  reference  to  the  ionic 
plane  of  cleavage.  Thus  we  have  incidentally  observed  many 
ways  in  which  acids,  bases,  and  salts  may  be  prepared,  that  do 
not  involve  a  union  of  the  constituent  ions  and  are  probably  not 
ionic. 

Oxygen  acids  can  almost  all  be  prepared  from  the  anhydride, 
that  is,  the  oxide  of  the  non-metal,  which  is  not  an  ionogen,  and 
water.  Phosphoric  acid,  sulphurous  acid  (p.  94),  hypochlorous 
acid  (C120  +  H20  — » 2HC10) ,  and  many  other  acids  are  so 
formed.  Hydrogen  fluoride,  chloride,  bromide,  and  iodide  are 
all  producible  by  union  of  the  constituent  elements.  Many  acids 
are  formed  from  others  when  the  latter  are  decomposed;  for 
example,  hydrochloric  acid  from  hypochlorous  acid  (p.  161). 

Bases  are  formed  by  the  union  of  oxides  of  metals  with  water 
(p.  94). 

The  dry  ways  of  forming  salts  are  very  numerous.  Thus,  many 
are  produced  by  direct  union  of  the  elements,  as  in  the  case  of  chlo- 
rides (p.  146),  sulphides  (p.  14),  and  other  simple  salts.  Many  are 
made  by  reduction  or  oxidation  from  other  salts,  as  potassium  chlo- 
ride from  potassium  chlorate  (p.  27),  or  potassium  perchlorate 
(q.v.)  from  the  latter.  Often  a  reducing  or  an  oxidizing  agent  is 
used,  as  in  making  sodium  nitrite  (see  index)  from  the  nitrate. 
Almost  all  oxygen  salts  can  be  obtained  by  the  union  of  two  oxides, 
as  calcium  carbonate  (see  index)  from  calcium  oxide  and  carbon 
dioxide.  Ammonium  salts  are  formed  by  combination  of  am- 
monia, which  is  not  an  ionogen,  with  acids  (p.  146). 

In  manufacturing  commercially  important  salts,  methods  like 
the  above,  as  well  as  those  involving  ionic  actions,  are  very  com- 
monly used.  In  each  case  the  cheapest  and  most  easily  acces- 
sible materials  are  chosen,  and  the  least  expensive  operation  is 
selected. 

Exercises.  —  1.  Give,  for  each  of  the  following,  a  definition, 
i.e.,  concise  description,  in  terms  of  experimental  facts:  acid  (pp.  52, 
158,  210,  246),  base  (pp.  94,  146,  246),  salt  (p.  246),  acid  salt, 
mixed  salt. 

2.  Give,  now,  a  definition  of  the  same  things  (see  1),  in  terms  of 
ions. 


262  COLLEGE    CHEMISTRY 

3.  Name  all  the  ionic  substances  whose  formulae  are  given  on 
pp.  212,  237,  and  classify  them  into  anions  and  cations. 

4.  Give  a  list  of  the  specific  physical  and  chemical  properties, 
including  those  that  can  be  used  as  tests,  of:  iodide-ion,  sulphate- 
ion,  cupric-ion,  chloride-ion. 

5.  Give  a  list  of  all  the  colorless  ionic  substances  you  can  think 
of. 

6.  Using  the  table  of  fractions  ionized  (p.  241),  prepare  lists  of 
the  pairs  of  ionic  substances  which  show  the  greatest,  and  the  least 
tendency  to  combine,  and  state  in  each  case  the  proportion  com- 
bining in  decinormal  solution. 

7.  In  the  case  of  the  green  solution  of  cupric  bromide  (p.  249), 
explain  in  detail  (p.  181)  the  effect  of  the  addition  of  potassium 
bromide.     Formulate  the  action  (p.  251).      j.; 

8.  In  the  case  of  the  chocolate-brown,  concentrated  solution  of 
cupric  bromide  (p.  249),  explain  in  detail  what  would  happen  to 
the  system:    (a)  if  metallic  zinc  were  to  be  added  (p.  259);   (6)  if 
hydrogen  sulphide  gas  were  to  be  led  into  the  solution  (CuS  is 
insoluble) . 

9.  Formulate,  after  the  models  on  pp.  251  and  252,  and  discuss 
fully,  the  interaction  of  ferric  chloride  and  ammonium  thiocyanate 
(p.  182). 

10.  What  is  implied  by  the  statements,  that  peroxides  are  salts 
and  that  hydrogen  peroxide  is  feebly  acid  (p.  223)? 

11.  Formulate  after  the  model  on  p.  252,  and  discuss  fully,  the 
interaction  of:   (a)  sodium  peroxide  and  hydrochloric  acid  (p.  222) ; 
(6)  barium  peroxide  and  sulphuric  acid. 

12.  Invent  an  interaction  of  two  soluble  salts  in  which  both 
products  shall  be  insoluble  (see  Table  of  Solubilities,  inside  of 
front  cover)  and  formulate  it,  (p.  252). 

13.  For  the  neutralization  of  77  c.c.  of  a  certain  alkaline  solution, 
25  c.c.  of  normal  hydrochloric  acid  are  required.     What  is  the 
normal  concentration  of  the  alkali?     If  the  alkali  was  sodium 
hydroxide,  what  weight  of  the  substance  was  present?     If  the 
alkali  was  barium  hydroxide,  what  weight  of  it  was  present? 

14.  Formulate  (p.  259)  the  actions  of  iron  and  of  aluminium  on 
dilute  hydrochloric  acid. 

15.  Formulate  (p.  259)  the  displacements  of  iodine  by  chlorine 
and  by  bromine  (p.  200). 


IONIC   SUBSTANCES  AND   THEIR  INTERACTIONS          263 

16.  Which  metals  (p.  260),  besides  platinum,  would  be  most 
likely  to  form  suitable  electrodes  for  an  electrolytic  cell? 

17.  To  which  classes  of  ionic  actions  do  those  of  iodine  on  hy- 
drogen sulphide  (p.  201),  and  of  calcium  on  cold  water  (p.  50), 
belong? 


CHAPTER  XX 
SULPHUR  AND  HYDROGEN  SULPHIDE 

Occurrence.  —  Free  sulphur  is  found  in  volcanic  regions  in 
Sicily,  where  it  is  mixed  with  gypsum  and  other  minerals  and  occu- 
pies the  pores  of  pumice-stone.  Rocky  materials  accompanying 
a  mineral  in  this  way  are  called  the  matrix.  The  other  important 
deposit  is  in  Louisiana.  There  are  many  minerals  containing 
sulphur  but,  with  the  exception  of  pyrite,  these  are  chiefly  impor- 
tant on  account  of  their  other  constituents.  Sulphides  of  metals, 
such  as  pyrite  FeSfe,  copper  pyrites  CuFeS2,  galena  PbS,  zinc- 
blende  ZnS,  and  sulphates,  like  gypsum  CaSO4,2H20,barite  BaS04, 
and  celestite  SrS04,  are  fairly  plentiful.  Sulphur  is  a  constituent 
of  the  proteins,  which  are  important  components  of  the  structure 
of  plants  and  animals. 

Manufacture.  —  In  Sicily,  sulphur  is  obtained  by  the  simple  ' 
process  of  melting  it  away  from  the  accompanying  volcanic  rock 
at  a  low  temperature.  The  liquid  sulphur  is  allowed  to  run  into 
wooden  molds,  in  which  it  solidifies  in  the  form  of  roll  sulphur,  or 
roll  brimstone.  To  produce  the  best  quality  it  is  subjected  to 
distillation  from  earthenware  retorts.  When  the  vapor  is  led  inljo 
a  large  brick  chamber,  it  condenses  upon  the  walls  and  floor  at 
first  in  the  form  of  flowers  of  sulphur,  and  later,  when  the 
chamber  becomes  heated,  as  a  liquid. 

In  Louisiana,  the  sulphur  forms  a  deposit  over  naif  a  mile  in 
diameter,  below  900  feet  of  clay,  quicksand,  and  rock.  It  is 
extracted  by  the  Frasch  method,  by  means  >  of  borings  which 
permit  four  pipes,  one  within  the  other,  to  reach  the  deposit. 
Water,  previously  heated  under  pressure  to  170°,  is  pumped  down 
the  two  outside  pipes  (6  and  8  inches  in  diameter).  After  time 
has  been  allowed  for  the  melting  of  a  mass  of  the  sulphur  (m.-p. 
114.5°),  compressed  air  is  forced  down  the  innermost,  one-inch 
pipe.  The  melted  sulphur  has  twice  the  specific  gravity  of  the 

264 


SULPHUR   AND   HYDROGEN   SULPHIDE  265 

water  in  the  outer  pipes.  But  the  mixture  of  air  and  sulphur  has 
about  the  same  specific  gravity,  and  so  flows  freely  up  the  three- 
inch  pipe  surrounding  the  air  pipe.  The  element  flows  into  a 
large,  wooden  enclosure,  in  which  it  solidifies,  and  is  practically 
pure  sulphur.  Each  well,  until  obstructed  by  collapse  of  the  rock 
and  quicksand  at  the  bottom,  produces  500  tons  a  day. 

The  greater  part  of  the  sulphur  of  commerce  formerly  came  from 
Sicily,  where,  in  1898,  447,000  tons  were  manufactured  against 
41,000  tons  elsewhere.  The  whole  supply  of  the  United  States 
(250,000  tons)  is  now  obtained  from  Louisiana.  The  world's 
consumption  is  over  800,000  tons. 

Physical  Properties.  —  The  chief  physical  peculiarity  of 
sulphur  is  that,  instead  of  appearing  in  only  three  familiar  physical 
states,  like  water,  it  possesses  two  familiar  and  perfectly  distinct 
solid  forms  and  two  different  liquid  states  of  aggregation. 

1.  Rhombic  Sulphur.     Native  sulphur  is  yellow,  has  a  sp.  gr. 
2.06  and  melts  at  112.8°.     It  is  almost  insoluble  in  water,  but 
dissolves  freely  in  carbon  disulphide  (41  parts  in  100  at  18°).     The 
crystals  of  native  sulphur,  as  well  as  those  obtained  by  evaporating 
a  solution,  belong  to  the  rhombic  system  (Fig.  7,  p.  12).     Roll 
sulphur  and  most  specimens  of  flowers  of  sulphur  are  the  same 
substance  although  the  crystals  in  their  growth  have  interfered 
with  one  another,  and  the  mass  is  crystalline,  simply,  and  not  well 
crystallized.     This  variety  is  called,  from  its  form,  rhombic  sul- 
phur.    This  form  is  stable  below  96°.     Above  that  temperature  it 
changes  slowly  into  monoclinic  sulphur. 

2.  Monoclinic  Sulphur.     When  a  large  mass  of 
melted  sulphur  solidifies  slowly,  and  the  crust  is 
pierced  and  the  remaining  liquid  poured  out  be- 
fore the  whole  has  become  solid,  the  interior  is  found 
to  be  lined  with  long,  transparent  needles  (Fig.  84). 
This  kind  of  sulphur  is  nearly  colorless,  has  a  sp.  gr. 

1.96,  melts  at  119.25°,  and  is  in  all  physical  re-         FIQ  ^ 
spects  a  different  individual  from  rhombic  sulphur. 
This  variety  is  named,  from  the  system  to  which  its  crystals 
belong,  monoclinic  sulphur.     This  form  can  be  kept  above  96° 
(transition  point,  p.  86),  but  when  allowed  to  cool,  it  slowly  be- 
comes opaque,  changing  into  particles  of  rhombic  sulphur. 


0 


266  COLLEGE    CHEMISTRY 

A  substance  which  has  two  solid  states  of  aggregation  and,  there- 
fore, two  crystalline  forms,  is  said  to  be  dimorphous  (two-formed). 

3.  S\  and  $M,  Vapor.  When  melted  sulphur  is  heated,  it  under- 
goes a  gradual  change,  which  is  especially  noticeable  near  160°. 
The  formerly  pale-yellow,  mobile  liquid  (S\)  suddenly  becomes 
dark-brown  in  color  and  so  viscous  (SM)  that  the  vessel  may  be 
inverted  without  loss  of  material:  S\  +±  SM.  The  liquid  is  a  mix- 
ture, containing  increasing  proportions  of  SM.  Beyond  260°  the 
viscidity  becomes  less,  and  at  444.7°  the  liquid  boils  and  passes  into 
sulphur  vapor. 

When  ordinary  sulphur  is  raised  to  the  boiling  point  and  then 
allowed  slowly  to  cool,  the  product  is  crystalline  and  soluble  in 
carbon  disulphide,  as  before.  The  change  from  Sx  to  SM  is  revers- 
ible. But  when  sulphur  is  boiled  and  then  suddenly  chilled  by 
pouring  into  cold  water,  it  is  at  first  semi-fluid.  After  several  days 
this  plastic  sulphur,  as  it  is  called,  becomes  hard.  It  is  then  found 
to  contain  rhombic  sulphur  mixed  with  30  per  cent  of  another 
variety  of  free  sulphur,  namely  SM.  This  part  is  almost  insoluble 
in  any  solvent.  Being  without  crystalline  structure,  it  is  called 
amorphous  (Gk.,  without  form)  sulphur.  Now  amorphous  bodies 
(see  Glass)  are  always  supercooled  liquids,  that  is,  liquids  still 
existing  as  such  at  a  temperature  at  which  the  solid,  crystalline 
form  is  the  stable  one.  This  is  simply  the  SM  in  a  supercooled 
state.  When  cold,  it  reverts  very  slowly  to  the  soluble  variety, 
and  years  are  required  for  the  completion  of  the  reversion  at 
room  temperature. 

Chemical  Properties.  —  At  low  temperatures  and  under  re- 
duced pressure,  the  formula  of  sulphur  vapor  is  SB.  As  the  tem- 
perature is  raised,  however,  the  vapor  expands  very  rapidly,  and 
at  800°  the  molecular  weight  is  64.2,  and  the  formula  therefore  S2 
(p.  117).  The  formula  of  dissolved  sulphur,  as  measured  by  the 
freezing-point  method  (p.  213),  is  Ss. 

Sulphur  is  an  active  chemical  substance  (p.  208).  When  finely 
divided  metals,  with  the  exception  of  gold  and  platinum  (pp.  60, 
260),  are  rubbed  together  with  powdered  sulphur,  union  takes 
place  and  sulphides  are  produced.  Sulphur  when  heated  com- 
bines with  great  vigor  with  iron  (p.  13),  copper,  and  most  of  the 
metals.  It  unites  also  with  many  of  the  non-metals.  Thus  with 


HYDROGEN   SULPHIDE  267 

oxygen  it  produces  sulphur  dioxide  (p.  31),  and  even  sulphur  tri- 
oxide  SOs.  It  unites  also  with  chlorine  directly.  When  sulphur 
is  treated  with  oxidizing  agents  in  presence  of  water,  no  trace  of 
sulphur  dioxide  (or  sulphurous  acid)  is  formed;  the  only  prod- 
uct is  sulphuric  acid  (see  p.  289).* 

Uses  of  Sulphur.  —  Large  quantities  of  crude  sulphur  are 
employed  for  making  sulphur  dioxide,  which  is  used  in  the  manu- 
facture of  sulphuric  acid,  in  bleaching  feathers,  straw,  and  wool, 
in  preserving  dried  fruits,  and  in  making  alkali  sulphites  for 
employment  in  the  bleaching  industry  and  in  paper-making.  The 
manufacture  of  carbon  disulphide  also  consumes  much  sulphur. 
Purified  sulphur  is  employed  in  the  manufacture  of  gunpowder, 
fireworks,  matches,  and,  by  combination  with  rubber,  of  vulcanite. 
Flowers  of  sulphur  is  used  in  vineyards  to  destroy  fungi,  which  it 
does  by  virtue  of  the  traces  of  sulphuric  acid  it  yields  by  oxidation. 

HYDKOGEN  SULPHIDE  H2S 

This  gas  is  found  dissolved  in  some  mineral  waters,  which  in  con- 
sequence are  known  as  sulphur  waters.  It  is  produced  in  the  de- 
composition of  animal  matter  containing  sulphur  (proteins),  when 
air  is  excluded.  Hence  the  odor  of  rotten  eggs  is  due  in  part  to  its 
presence. 

Preparation.  —  1.  Hydrogen  and  sulphur  do  not  unite  percep- 
tibly in  the  cold.  At  310°  almost  complete  union  occurs,  but  about 
168  hours  are  required  for  the  attainment  of  equilibrium. 

2.  Sulphides  of  metals,  being  salts,  are  acted  upon  more  or  less 
easily  by  dilute  acids,  and  give  hydrogen  sulphide.  Ferrous  sul- 
phide, the  least  expensive  of  those  easily  affected,  is  generally 
used: 

FeS  +  2  HC1  fc*  H2S  t  +  FeCl2. 

For  hydrochloric  acid  we  may  substitute  an  aqueous  solution  of 
any  active,  non-oxidizing  acid  (see  p.  268,  last  line).  A  Kipp's 
apparatus  (p.  54)  is  commonly  employed. 

*  The  paragraph  on  the  chemical  relations  of  the  element  (see  end  of  this 
chapter)  should  be  read  at  this  point. 


268  COLLEGE    CHEMISTRY 

3.  Hydrogen  sulphide  is  the  invariable  product  of  the  extreme 
reduction  of  any  sulphur  compound.  Thus,  it  is  formed  by  the 
actton  of  hydrogen  iodide  upon  concentrated  sulphuric  acid  (p. 
201).  Even  sulphur  itself  is  reduced  by  dry,  gaseous  hydrogen 
iodide  : 


Physical  Properties.  —  Hydrogen  sulphide  is  a  colorless  gas 
with  a  characteristic  odor.  When  liquefied,  it  boils  afc  —62°,  and 
in  solid  form  melts  at  -83°.  The  solubility  in  water  at  10°  is  360 
volumes  in  100,  and  becomes  less  as  the  temperature  is  raised. 
The  gas  can  be  driven  out  completely  by  boiling  the  solution  (c/. 
p.  145).  The  gas  is  very  poisonous,  one  part  in  two  hundred  of 
air  being  fatal  to  mammals. 

Chemical  Properties  of  Hydrogen  Sulphide  Gas.  —  When 
heated,  the  gas  dissociates: 


At  310°  the  decomposition  is  slight  (c/.  p.  267),  but  becomes 
greater  at  higher  temperatures. 

The  gas  burns  in  air,  forming  steam  and  sulphur  dioxide.  The 
temperature  of  the  mantle  of  flame  surrounding  the  gas,  as  it  issues 
from  a  jet,  being  far  above  310°,  the  gas  in  the  interior  is  dissociated 
before  it  meets  with  any  oxygen.  Hence  a 
cold  dish  held  across  the  flame  (Fig.  85)  re- 
ceives a  deposit  of  free  sulphur,  and  a  part  of 
the  hydrogen  also  escapes  unburnt.  It  may 
be  remarked  that  dissociation  of  this  kind 
probably  precedes  the  combustion  of  most 
gaseous  compounds  (see  Flame). 

The  metals,  down  to  and  including  silver 
FIG  85  m  ^ne  electromotive   series,   when   exposed 

to  the  gas,  quickly  receive  a  coating  of  sul- 
phide. The  tarnishing  of  silver  in  the  household  is  probably  due 
to  a  trace  of  hydrogen  sulphide  in  the  illuminating  gas  which 
escapes  from  slight  leaks  in  the  pipes.  That  the  gas  should  thus 
behave  like  free  sulphur  shows  its  instability. 

This  instability  is  shown  also  in  the  fact  that  it  reduces  sub- 


HYDROGEN   SULPHIDE  269 

stances,  such  as  sulphur  dioxide,  which  are  not  affected  by  free 
hydrogen : 

S02  +  2H2S-»2H2O  +  3S. 

This  action  takes  place  in  the  cold,  and  much  more  rapidly  when 
the  gases  are  moist  than  when  they  are  dry  (p.  160).  Native  sul- 
phur is  often  produced  by  this  action,  as  both  of  these  gases  are 
found  issuing  from  the  ground  in  volcanic  neighborhoods.  Sul- 
phur is  deposited  also  when  hydrogen  sulphide  undergoes  a  partial 
combustion  with  a  restricted  supply  of  oxygen,  2H2S  +  O2  — > 
2H2O  +  2S,  and  its  formation  in  nature  is  sometimes  to  be  ac- 
counted for  in  this  way. 

A   Characteristic  of  Reduction  and  Oxidation.  —  In  the 

former  of  the  two  actions  last  mentioned,  it  will  be  seen  that,  while 
the  S02  was  reduced  to  S,  at  the  same  time  H2S  was  oxidized  (to  S). 
In  the  second  action,  H2S  was  oxidized  to  S,  and  02  was  reduced  to 
2H20.  It  is  a  characteristic  of  such  actions  that  one  substance  is 
oxidized  and  another  reduced:  oxidation  and  reduction  always 
occur  together,  in  the  same  reaction.  Here,  under  hydrogen 
sulphide,  we  speak  of  its  reducing  effect  on  sulphur  dioxide.  Under 
sulphur  dioxide,  however,  we  should  speak  of  the  oxidizing  effect 
of  the  substance  on  hydrogen  sulphide. 

Chemical  Properties  of  the  Aqueous  Solution  of  Hydrogen 
Sulphide.  —  While  the  gas  itself  is  not  an  acid,  its  solution  in 
water  gives  a  feeble  acid  reaction  with  litmus,  and  is  sometimes 
named  hydrosulphuric  acid  H2S,  Aq.  The  conductivity  of  a  N/1Q 
aqueous  solution  is  small,  and  only  0.0007  (0.07  per  cent)  of  the 
substance  is  ionized: 

H2S  ±3  H+  +  HS~  («=*  H+  +  S=). 

Some  S=  ions  are  present.  But  hydrosulphide-ion  HS~,  although 
an  acid,  is  less  dissociated  than  is  water  itself,  and  the  amount  of 
sulphide-ion  is  therefore  very  small.  The  salts  of  hydrosulphide- 
ion,  such  as  NaHS  (sodium  acid  sulphide,  see  next  section),  give 
therefore  neutral  solutions.  This  behavior  is  the  rule  with  the 
acid  salts  of  feeble  dibasic  acids  (p.  241). 

As  an  acid,  the  solution  of  hydrogen  sulphide  may  be  neutralized 

Z4t«^S4-Ow      -     2-  #Vl>    4       L$> 

"      4  .r 


270  COLLEGE    CHEMISTRY 

by  bases.  For  the  same  reason  it  enters  into  double  decomposition 
with  salts  (see  next  section). 

By  the  action  of  oxygen  from  the  air  upon  an  aqueous  solution  of 
hydrogen  sulphide,  the  sulphur  is  slowly  displaced  and  appears  in 
the  form  of  a  fine  white  powder: 

02  +  2H2S  -»  2S I  +  2H20. 

This  is  an  action  similar  to  the  displacement  of  ionic  bromine  by 
free  chlorine  (p.  259). 

The  solution  of  the  gas  is  a  reducing  agent,  as  its  action  upon 
iodine  shows  (p.  202) .  So,  also,  in  presence  of  an  acid,  it  removes 
oxygen  from  dichromic  acid  (produced  by  the  action  of  an  acid 
upon  potassium  dichromate) : 

K2Cr2O7  +  2HC1  +±  H2Cr207  +  2KC1.  (1) 

H2Cr207  +  6HC1  -»  4H20  +  2CrCl3  ( +  30) .      (2) 

(3O)  +  3H2S  ->  3H2O  +  3S. (3) 

Adding:  K2Cr207  +  8HC1  +  3H2S  -» 2KC1  +  2CrCl3  +  7H2O  +  38. 

The  first  partial  equation  (cf.  p.  194)  represents  the  regular  inter- 
action of  two  ionogens,  but  the  second  interaction  does  not  take 
place  unless  an  oxidizable  body  (here  the  hydrogen  sulphide)  is 
present  to  take  possession  of  the  oxygen  which  it  is  capable  of 
delivering  (cf.  p.  225). 

The  foregoing  illustrates  a  fourth  kind  of  ionic  chemical  change 
(p.  259),  namely  that  in  which  a  compound  ion  is  formed  or  decom- 
posed. Here  dichromate-ion  Cr2O7—  gives  chromic-ion  Cr+++  and 
water.  For  other  illustrations  see  pp.  56,  161,  206,  224,  225,  274. 

Sulphides.  —  As  a  dibasic  acid  (p.  269),  hydrogen  sulphide 
gives  both  acid  and  normal  (or  " neutral")  sulphides,  such  as 
NaHS  and  Na^S. 

The  acid  sulphides  are  obtained  by  passing  the  gas  in  excess  into 
solutions  of  soluble  bases: 

H2S  +  NaOH  -»  H20  +  NaHS, 

and  are  neutral  in  reaction.  Their  negative  ion,  HS7,  is  not  further 
dissociated  (see  preceding  section). 

By  adding  to  the  above-mentioned  solution  an  amount  of  sodium 
hydroxide  equal  to  that  used  before,  and  driving  off  the  water  by 


HYDROGEN   SULPHIDE  271 

evaporation,  the  second  unit  of  hydrogen  is  displaced,  and  nor- 
mal ("neutral")  sodium  sulphide  is  formed: 

NaOH>  NaHS  fc;  Na2S  +  H20  1  . 

This  action  is  wholly  reversed  when  the  dry  sodium  sulphide  is 
dissolved  in  water,  the  salt  being  completely  hydrolyzed  (p.  197)  to 
the  acid  salt: 


H20±=»OET  +  H+J  " 

The  HS~  gives  a  lower  concentration  of  hydrogen-ion  than  the 
water,  and  hence  uses  up  in  its  formation  the  ions  of  hydrogen 
produced  by  the  latter,  until  an  amount  of  hydroxide-ion  equiva- 
lent to  half  the  sodium  is  formed.  The  abbreviated  equation 
shows  this  more  clearly: 

S=  +  H+  +  OH~  ->  HS~  +  OH". 

The  solution  is  therefore  strongly  alkaline  in  reaction.  In  general, 
a  normal  salt  derived  from,  an  active  base  and  a  weak  acid  is  hydro- 
lyzed to  sorne  extent  by  water  and  gives  an  alkaline  solution. 

In  the  abbreviated  formulation  used  above,  the  union  of  Na+ 
and  OH~  to  form  NaOH  is  not  shown,  because  it  is  slight  in  dilute 
solution  and  does  not  affect  the  result.  The  union  'of  S=  and  H+ 
to  form  HS~  is  alone  shown,  because  it  is  extensive  and  significant. 
To  save  space,  this  plan  will  be  used  in  future,  where  the  same 
situation  exists. 

The  soluble  acid  sulphides  are  oxidized  in  aqueous  solution  by 
atmospheric  oxygen: 

2NaSH  +  02  ->  2NaOH  +  2S. 

The  sulphur  is  not  precipitated,  but  combines  with  the  excess  of  the 
sulphide,  forming  polysulphides  (see  below).  Some  sodium  thio- 
sulphate  is  produced  at  the  same  time. 

The  Action  of  Acids  on  Insoluble  Sulphides.  —  The  inter- 
action of  sulphides  and  acids  is  itself  so  important  a  matter  in 
chemistry,  and  is  so  similar  in  theory  to  many  other  kinds  of 
actions,  that  special  attention  should  be  given  to  it.  The  common 
method  of  preparing  hydrogen  sulphide  from  ferrous  sulphide 
affords  a  suitable  illustration. 


I 


/  I ', 


272  COLLEGE    CHEMISTRY 

Since  ferrous  sulphide  is  but  slightly  soluble  in  water,  the  action 
proceeds  by  a  rather  complex  series  of  equilibria: 

FeS  (solid)  ±+FeS  (dslvd)  fc?  Fe++  +  S= 


It  will  be  seen  that  a  number  of  reversible  changes  are  involved, 
and  the  question  is,  -why  does  the  reaction  proceed  forward,  as  it 
does?  To  answer  this  question,  a  consideration  of  each  of  the 
equilibria,  separately,  is  required. 

1.  The  dissolved  hydrogen  sulphide  is  very  feebly  ionized,  and 
maintains  a  smaller  concentration  of  sulphide-ion  S—  than  does 
ferrous  sulphide,  in  spite  of  the  comparative  insolubility  of  the 
latter.     Hence,  the  S=  formed  from  the  FeS  is  continuously  re- 
moved by  union  with  the  hydrogen-ion  furnished  by  the  acid, 
S=  +  2H+  t^  H2S,  and  all  the  other  equilibria  are  constantly  dis- 
placed forward  on  this  account.      The  action  is  therefore,  in 
essence,  like  neutralization  (p.  254). 

2.  The  union  of  S=  and  2H+  depends  on  the  magnitude  of  the 
-product  of  their  concentrations  (p.  184),  [S=]  X  [H+]  X  [H+],  or 
[S=]  X  [H+]2.     Hence,  although  [S=]  is  minute,  on  account  of  the 
insolubility  of  FeS,  [H+]  is  large  on  account  of  the  great  dissocia- 
tion of  the  HC1  and  the  fact  that  a  strong  solution  of  the  acid  can 
be  used.     Thus  the  product  may  be  large  enough  for  the  purpose. 

3.  When  a  still  more  insoluble  sulphide,  like  cupric  sulphide 
CuS  is  employed,  the  concentration  of  the  sulphide-ion  [S=]  is  too 
small  to  play  its  part  and  the  action  makes  almost  no  progress.    In 
this  case,  a  concentration  of  H+,  sufficient  to  raise  the  product  to 
the  necessary  value,  cannot  be  obtained  with  any  acid. 

4.  The  fact  that  hydrogen  sulphide  is  fairly  soluble  (3.6  vols.  :  1 
vol.)  hinders  the  action.     It  prevents  that  free  escape  of  one  prod- 
uct which  is  so  constantly  a  factor  in  promoting  reversible  chemical 
changes.     Thus,  if  cadmium  sulphide  CdS,  which  lies  between 
ferrous  and  cupric  sulphides,  in  solubility,  is  employed  along  with 
rather  dilute  hydrochloric  acid,  a  concentration  of  hydrogen  sul- 
phide sufficient  to  stop  the  action  accumulates  before  the  liquid  is 
saturated  with  the  gas,  and  the  latter  can  begin  to  escape.     There 
are  then  two  ways  of  making  this  action  continuous.     Either 
stronger  hydrochloric  acid,  giving  a  higher  concentration  of  H+ 
may  be  used  to  force  the  formation  of  more  H2S  (by  union  of  2H+ 


HYDROGEN   SULPHIDE  273 

and  S=),  or  the  reverse  action,  due  to  accumulation  of  H2S  (dslvd), 
may  be  diminished  mechanically  by  leading  air  through  the  mix- 
ture (p.  129)  and  so  removing  the  hydrogen  sulphide  as  fast  as  it 
is  formed.  Either  plan  will  cause  complete  interaction  with  the 
cadmium  sulphide. 

Classification  of  Insoluble  Sulphides.  —  In  analytical 
chemistry,  advantage  is  taken  of  the  different  solubilities  of  the 
sulphides,  for  the  purpose  of  identifying  the  metallic  elements,  and 
of  separating  mixtures  containing  several  such  elements.  Three 
classes  are  distinguished. 

1.  The  sulphides  of  silver,  copper,  mercury,  and  some  other 
metals  are  exceedingly  insoluble,  and,  therefore,  do  not  interact 
with  dilute  acids  as  does  ferrous  sulphide  (p.  271).     These  may 
therefore  be  made  by  leading  hydrogen  sulphide  into  solutions  of 
their  salts: 

CuS04  +  H2S  t?  CuS  |  +  H2S04. 

The  acid  produced  has  scarcely  any  effect  upon  the  sulphide,  and 
almost  no  reverse  action  is  observed.  In  this  action  the  sulphide- 
ion  is  the  active  substance  and,  by  its  removal,  all  the  equilibria 
are  displaced  forwards. 

2.  The  sulphides  of  iron,  zinc,  and  certain  other  metals  are  insol- 
uble in  water,  but  not  so  much  so  as  the  last  class.     Hence  they  are 
decomposed  by  dilute  acids,  and  the  reverse  of  the  above  action 
takes  place  almost  completely.     These  sulphides  must  therefore  be 
made,  either  by  combination  of  the  elements,  or  by  adding  a  soluble 
sulphide  to  a  solution  of  a  salt: 

FeS04  +  (NH4)2S^FeS|+  (NH4)2S04. 

No  acid  is  produced  in  this  sort  of  interaction,  and  the  considerable 
insolubility  of  the  sulphide  of  iron  or  zinc  in  water  renders  the 
change  nearly  complete. 

3.  The  sulphides  of  barium,  calcium,  and  some  other  metals 
(q.v.),  although  insoluble  in  water,  are  hydrolyzed  by  it,  and  give 
soluble  products,  the  hydroxide  and  hydrosulphide : 

2CaS  +  2H2O  fc*  Ca(OH)2  +  Ca(SH)2. 

They  may  be  prepared  by  direct  union  of  the  elements,  and  from 
the  sulphates  by  reduction  with  carbon.  But  they  are  not  pre- 
cipitated by  hyolrogen  sulphide  or  ammonium  sulphide. 


u  et1 

274  COLLEGE    CHEMISTRY 

Poly  sulphides.  —  When  sulphur  is  shaken  with  a  solution  of  a 
soluble  sulphide  or  acid  sulphide,  such  as  sodium  sulphide,  it  dis- 
solves, and  evaporation  of  the  solution  leaves  residues,  varying  in 
composition  from  Na^  to  Na^Ss.  These  appear  to  be  mixtures 
composed  mainly  of  NagS  and  Na^. 

When  an  acid  is  poured  into  sodium  polysulphide  solution, 
minute  spherules  of  rhombic  sulphur  are  precipitated: 

2HC1  ->  2NaCl  +  H2S  |  +  3S  [  . 


The  Chemical  Relations  of  the  Element  Sulphur.  —  In 

combination  with  metals  and  hydrogen,  sulphur  is  bivalent,  form- 
ing compounds  like  H2S,  FeS,  CuS,  and  HgS.  In  combination  with 
non-metals,  however,  the  valence  is  frequently  greater,  the  maxi- 
mum being  seen  in  sulphur  trioxide,  where  the  sulphur  is  sexivalent. 
Its  oxides  are  acid-forming,  and  it  is,  therefore,  a  non-metal. 

Exercises.  —  1.  How  could  the  decomposition  of  hydrogen  sul- 
phide at  310°  be  rendered,  (a)  more  complete,  (6)  less  complete? 
Would  the  percentage  decomposed  be  affected,  (a)  by  reducing  the 
pressure,  (6)  by  mixing  the  gas  with  an  indifferent  gas? 

2.  What  are  the  relative  volumes  of  the  gases  (p.  150)  in  the 
action  of,  (a)  hydrogen  iodide  and  sulphur,  (6)  hydrogen  sulphide 
and  sulphur  dioxide?         J   tf  ^  ~S  ^    //  z;  y  \~  j  '* 

3.  To  what  classes  of  ionicSictions  (p.  §59)  0:0  the  interactions  of 
hydrogen  sulphide  solution  with,  (a)  oxygen  (p.^270),  (6)  sodium 
hydroxide  (p.  270),  (c)  iodine  (p.  202)  belong? 

4.  Show  which  actions  on  the  pages  referred  to  on  p.  270  illus- 
trate the  fourth  kind  of  ionic  chemical  change,  and  how  they  do  so? 

5.  Why  is  normal  sodium  sulphide  only  half  hydrolyzed  by 
water? 

6.  Formulate  completely,  after  the  model  on  p.  252,  the  actions 
of  (a)  hydrogen  sulphide  and  cupric  sulphate  solution;    (6)  am- 
monium sulphide  and  ferrous  sulphate.     In  each  case  explain 
which  equilibrium  determines  the  direction  of  the  action. 


1    -f  7  -> 


CHAPTER  XXI 
THE   OXIDES  AND   OXYGEN  ACIDS   OF   SULPHUR 

THE  only  important  oxides  of  sulphur  are  the  dioxide  S02  and 
the  trioxide  SOs.  They  are  the  anhydrides  (p.  94)  of  sulphurous 
acid  H2SOs  and  of  sulphuric  acid  H2S04,  respectively. 

The  Preparation  of  Sulphur  Dioxide  SO2.  —  1.  When  sulphur 
burns  in  air  or  oxygen,  sulphur  dioxide  is  produced  (p.  32).  2. 
The  larger  part  of  the  sulphur  dioxide  used  in  commerce  is  probably 
obtained  by  the  roasting  (calcining)  of  sulphur  ores.  Pyrite  FeS2, 
for  example,  on  account  of  the  large  amount  of  sulphur  which  it 
contains,  can  be  burnt  in  a  suitable  furnace: 

4FeS2  +  1102  ->  2Fe2O3  +  8S02  f . 

The  gas,  although  mixed  with  great  amount  of  nitrogen  which 
entered  as  part  of  the  air,  can  be  used  to  make  sulphuric  acid. 

It  should  be  noted,  in  passing,  that  heating  and  roasting  or  cal- 
cining are  distinct  processes  in  chemistry.  Roasting  or  calcining 
always  assumes  the  access  of  the  air  and  employment  of  its  oxygen; 
heating,  in  the  absence  of  modifying  words,  assumes  the  exclusion 
or  the  chemical  indifference  of  the  air. 

3.  In  the  laboratory,  a  steady  stream  of  the  gas  is  obtained  by 
allowing  hydrochloric  acid  to  drop  upon  solid  sodium  acid  sul- 
phite, or  concentrated  sulphuric  acid  to  trickle  into  a  40  per  cent 
solution  of  the  same  salt  (Fig.  24,  p.  54) : 

HC1  +  NaHS03  fc?  NaCl  +  H2S03  fc?  H2O  +  SQjt- 

The  sulphurous  acid,  being  very  unstable,  decomposes  spontane- 
ously into  water  and  sulphur  dioxide,  and  the  latter  escapes  when 
sufficient  water  for  its  solution  is  not  present. 

4.  Sulphur  dioxide  can  also  be  made  by  the  reduction  of  con- 
centrated sulphuric  acid  by  copper  at  a  high  temperature.     A  part 

275 


276  COLLEGE    CHEMISTRY 

of  the  acid 'loses  oxygen  to  form  water  with  the  hydrogen  of 

another  molecule: 

Partial  I :  H2S04  -»  H20  +  SO2  (+  0). 

Partial  2:     (0)  +  H2SO4  +  Cu  -»  H2O  +  CuSO4. 

2H2S04  +  Cu  ->  2H2O  +  S02  +  CuS04. 

Some  easily  oxjplized  non-metals,  such  as  carbon  and  sulphur,  act 
in  the  same  way,  C  +  2H2S04  -» 2H20  +  2S02  +  C02. 

Making  Equations  by  Positive  and  Negative  Valences.  — 

Equations  like  the  foregoing  can  be  constructed  also  by  assuming 
that  each  element  in  a  compound  is  either  positive  or  negative, 
and  by  marking  the  valences  accordingly  (for  details,  see  p.  322). 
Thus,  in  sulphuric  acid,  we  have  2H+  (positive,  univalent)  and  40= 
(each  bivalent  and  negative).  Since  the  numbers  of  positive  and 
negative  valences  must  be  equal,  and  we  have  2©*  and  8©,  it 
follows  that  the  sulphur  carries  6©,  Sttt. 

Now  when,  in  making  the  experiment,  we  find  the  products  S02 
and  CuS04,  we  may  infer  that  the  hydrogen  formed  water.  We 
infer,  also,  that  to  obtain  two  compounds  containing  sulphur,  at 
least  2H2S04  was  required.  We  then  note  that  the  S  in  S02  is 
quadrivalent.  Hence  Sttt  became  Stt  and  2©  were  released. 
The  metallic  copper  used  was  free  and  without  valence,  and  be- 
came CuSO4,  in  which  it  is  GU++.  It  obtained  the  2©  from  the 
sulphur.  The  action  can  therefore  be  analyzed  as  follows: 

[2H+  +  Sttt  +  40=]  -»  [Stt +  20=]+  [2H+  +  0=]  +  [0=  +  2©] 
First  H2S04  S02  H2O  Balance 

The  second  H2S04  gives  [2H+  +  S04=].  The  Cu  takes  the  2© 
giving  Cu++,  and  this  with  the  S04=  gives  CuS04.  The  2H+  takes 
the  0=  from  the  balance,  giving  H2O.  Thus,  the  whole  balance  is 
used  and  the  products  are  accounted  for.  The  equation  must 
therefore  be: 

2H2S04  +  Cu  ->  S02  +  2H20  +  CuS04. 

It  will  be  noted  that  the  two  molecules  of  sulphuric  acid  play 
different  roles.  Only  one  of  them  is  used  in  oxidizing. 

*  The  signs  ©  and  ©  stand  for  quantities  of  electricity  equal  to  those 
carried  by  one  equivalent  of  an  ionic  substance,  and  therefore  required  for 
its  discharge  and  liberation. 


THE    OXIDES   AND    OXYGEN  ACIDS    OF   SULPHUR         277 


Similarly,  with  sulphuric  acid  and  carbon,  the  same  analyzed 
equation  applies.  The  carbon  gives  CO2.  Thus,  the  carbon  goes 
from  C°  to  Ctt.  To  obtain  the  40,  2H2SO4  is  required  (equation 
above).  Hence, 

2H2S04  +  C  ->  C02  +  2S02  +  2H20. 

When  hydrogen  sulphide  is  led  through  concentrated  sulphuric 
acid,  the  latter  is  reduced  to  sulphur  dioxide,  and  the  former  is 
oxidized,  giving  free  sulphur  (p.  270) : 

2H+  +  S=  +  20  ->2H+  +  S°J. 

Since  this  action  requires  20,  and  sulphuric  acid  in  giving  S02 
delivers  20,  it  follows  that  1H2SO4  will  decompose  1H2S: 
H2S04  +  H2S  -»  2H20  +  S°  j,  +  S02. 

Finally,  when  HI  with  sulphuric  acid  (p.  201)  gives  free  iodine 
(1°),  and  H2S  (2H+  +  S=),  evidently  Sffi  in  sulphuric  acid  gives 
up  80,  becoming  S~: 

[2H+  +  Sttt  +  40=]  ->  2H+  +  S=  +  4O=  +  80 
and  [H+  +  I~]  +  0  ->  H+  +  P. 

Evidently,  1H2SO4  giving  80  will  interact  with  SHI,  changing 
8I~  into  81°.     Hence, 

H2S04  +  SHI  -»  4H20  +  H2S  +  81°. 

The  reader  should  practice  the  use  of  this  method  by  making  the 
equations  for  the  actions  of  zinc  (p.  268  giving 
hydrogen  sulphide)  and  of  hydrogen  bromide 
(p.  196)  upon  sulphuric  acid. 

Physical  and   Chemical  Properties.  — 

Sulphur  dioxide  is  a  gas  possessing  a  pene- 
trating and  characteristic  odor.  This  is  fre- 
quently spoken  of  as  the  "odor  of  sulphur," 
but  it  should  be  remembered  that  sulphur 
itself  has  scarcely  any  smell  at  all.  The 
weight  of  the  G.M.V.  of  the  gas  (65.54  g.)  FJQ 

shows  it  to  be  more  than  twice  as  heavy  as  air. 
By  means  of  a  freezing  mixture  of  ice  and  salt  (Fig.  86),  the  gas 
is  easily  condensed  in  a  U-tube  to  a  transparent  mobile  fluid,  which 


278  COLLEGE    CHEMISTRY 

boils  at  —8°.  At  20°,  the  liquid  gives  a  vapor  pressure  of  only 
3|  atmospheres,  so  that  the  liquid  is  handled  and  sold  in  glass 
syphons  or  in  sealed  tin  cans.  The  solubility  of  the  'gas  in  water 
is  5000  volumes  in  100.  The  liquid  is  completely  freed  from  the 
gas  by  boiling  (cf.  p.  145). 

As  regards  chemical  properties,  sulphur  dioxide  is  stable  (p.  93). 

It  unites  with  water  to  form  sulphurous  acid  H2S03,  which  is 
unstable,  and  exists  only  in  solution. 

Since  the  maximum  valence  of  sulphur  is  6,  sulphur  dioxide,  in 
which  but  four  of  the  valences  of  sulphur  are  used,  is  unsaturated. 
It  is  therefore  still  able  to  combine  directly  with  suitable  elements, 
such  as  chlorine  and  oxygen.  When  it  is  mixed  with  chlorine  in 
sunlight,  a  liquid,  sulphuryl  chloride  SO2C12  is  produced. 

Liquefied  sulphur  dioxide  is  employed  for  bleaching  straw,  wool, 
and  silk  (see  p.  289).  As  a  disinfectant  it  has  been  displaced  to  a 
large  extent  by  formaldehyde. 

The  Liquefiability  of  Gases.  —  It  will  assist  us  in  recalling 
which  gases  are  hard  to  liquefy  and  which  easy,  if  we  memorize 
the  fact  that  Faraday  (from  1823  to  1845)  liquefied  most  of  the 
familiar  gases  and  failed  only  with  three,  namely  hydrogen  (c.t. 
-242°),  oxygen  (c.t.  -113°),  and  nitrogen  (c.t.  -146°).  These, 
with  nitric  oxide  NO  (c.t.  -93.5°),  carbon  monoxide  CO  (c.t. -40°), 
methane  CEU  (c.t.  -99°),  and  the  six  inert  gases  (pp.  335-337), 
are  the  ones  which  have  low  critical  temperatures  (cf.  p.  78)  and 
are  difficult  to  liquefy. 

Of  the  gases  we  have  studied,  the  ones  which  are  more  or  less 
easily  liquefied  are:  hydrogen  chloride  (c.t.  +52°),  bromide,  and 
iodide,  chlorine  (c.t.  +141°),  ozone,  hydrogen  sulphide  (c.t.  +  100°), 
sulphur  dioxide  (c.t.  +154°). 

The  Solubilities  of  Gases.  —  For  the  purpose  of  remember- 
ing the  solubilities  of  gases  in  water,  it  is  convenient  to  divide  the 
gases  into  three  classes.  The  following  are  the  ones  we  have 
studied: 

1.  Slightly   soluble:     Oxygen    (4   vol.  :  100   at   0°),   hydrogen 
(2  :  100  at  0°). 

2.  Soluble:    Chlorine    (260  vol.  :  100  at   10°),   hydrogen  sul- 
phide (440  :  100  at  0°). 


THE    OXIDES   AND    OXYGEN   ACIDS    OF   SULPHUR         279 

3.  Very  soluble:  Hydrogen  chloride  (505  vol.  :  1  at  0°),  bro- 
mide (404  :  1)  and  iodide  (1570  :  1),  sulphur  dioxide  (69  :  1  atO°). 

Preparation  of  Sulphur  Trioxide  S03.  —  Although  the  for- 
mation of  sulphur  trioxide  is  accompanied  by  the  liberation  of  much 
heat,  sulphur  dioxide  and  oxygen,  whether  cold  or  warm,  unite 
very  slowly.  Ozone,  however,  combines  with  the  former  readily. 

The  interaction  of  sulphur  dioxide  and  oxygen  is  hastened  by 
finely  divided  platinum,  which  remains  itself  unchanged  and  simply 
acts  as  a  catalytic  agent.  The  contact  "process,  as  this  is  called, 
has  been  rendered  available  for  the  commercial  manufacture  of 
sulphur  trioxide  by  Knietsch  (1901).  At  400°,  the  temperature 
used,  98-99  per  cent  of  the  materials  unite. 

02  +  2S02  ->  2S03  +  2  X  22,600  cal. 

Below  400°,  the  union  is  too  slow.  Above  400°,  the  reverse  action 
is  strengthened  (Van't  HofFs  law,  p.  188),  and  the  union  is  too 
incomplete.  The  vaporous  product  is  condensed  by  being  led 
into  97-99  per  cent  sulphuric  acid,  and  the  concentration  of  the 
liquid  is  constantly  maintained  at  this  point  by  the  regulated  in- 
flux of  water.  The  sulphur  dioxide  is  obtained  by  calcining  ores 
(p.  275).  These  contain  impurities  which  must  be  removed  very 
thoroughly.  Dust  from  the  roasting  and  oxide  of  arsenic,  which 
are  present,  will  otherwise  "poison"  the  contact  agent  (platinum 
or  ferric  oxide)  and  soon  almost  stop  the  union. 

The  process  may  be  illustrated  by  placing  some  platinized 
asbestos*  in  a  tube  (Fig.  66,  p.  156),  which  is  gently  warmed,  and 
introducing  oxygen  and  sulphur  dioxide  through  the  limbs  of  the 
Y-tube.  Dense  fumes  appear  at  the  exit  (see  next  section). 

Formerly  sulphur  trioxide  was  obtained  by  the  distillation  of 
impure  ferric  sulphate,  Fe2(S04)3  — >  Fe203  +  3S03. 

Physical  and  Chemical  Properties.  —  Sulphur  trioxide  is  a 
volatile  liquid  (b.-p.  46°).  The  crystals,  obtained  by  cooling,  melt 
at  14.8°.  It  fumes  strongly  when  exposed  to  the  air,  in  conse- 
quence of  the  union  of  the  vapor  with  moisture  and  the  production 
of  minute  drops  of  sulphuric  acid.  A  white  crystalline  variety, 

*  Asbestos,  dipped  in  a  solution  of  chloroplatinic  acid  and  heated  in  the 
Bunsen  flame:  H2PtCl6  -»  Pt  +  2HC1 1  +  2C12 1 . 


280  COLLEGE    CHEMISTRY 

closely  resembling  asbestos  in  appearance,  is  the  more  familiar 
form  of  the  substance,  which  is  dimorphous  (p.  266). 

As  to  chemical  properties,  the  vapor  of  sulphur  trioxide  dissociates 
into  sulphur  dioxide  and  oxygen  (400°,  2%;  700°,  40%). 

Sulphur  trioxide  is  not  itself  an  acid,  but  it  is  the  anhydride 
of  sulphuric  acid.  When  placed  in  water  it  unites  vigorously, 
causing  a  hissing  noise  due  to  the  steam  produced  by  the  heat  of 
the  union. 

Just  as  sulphur  trioxide  unites  with  water  to  give  hydrogen 
sulphate,  so  it  combines  vigorously  with  many  oxides  of  metals, 
producing  the  corresponding  sulphates : 

H20  +  S03  fc?  H2SO4,  CaO  +  S03  ->  CaS04. 

The  union  of  an  oxide  of  a  non-metal  with  the  oxide  of  a  metal,  in 
this  fashion,  is  a  general  method  of  obtaining  salts  (cf.  p.  261). 

Oxygen  Acids  of  Sulphur.  —  Sulphurous  and  sulphuric  acids 
have  been  mentioned  frequently  already.  Next  to  them  in  im- 
portance come  thiosulphuric  acid  and  persulphuric  acid.  The 
compositions  of  the  acids  show  their  relationships: 

Hyposulphurows  acid,  H^C^.  Sodium  hyposulphite,  Na2S2O4. 

Sulphurous  acid,  H2SO3.  Sodium  sulphite,  Na2SO3. 

Sulphuric  acid,  H2SC>4.  Sodium  sulphate,  Na-zSCX 

Thiosulphuric  acid,       H^Oa.  Sodium  thiosulphate,  Na2S208. 

Persulphuric  acid,         H2S2O8.  Sodium  persulphate,  Na2S2O8. 

Thiosulphuric  acid  (Gk.  0etov,  sulphur)  is  so  named  because  it 
contains  one  unit  of  sulphur  in  place  of  one  of  the  units  of  oxygen 
of  sulphuric  acid.  Note  that  when  the  names  of  the  acids  end  in 
ous  and  ic,  the  names  of  the  salts  end  in  ite  and  ate,  respectively. 
Besides  the  above  we  have  also  the  polythionic  acids,  namely: 
dithionic  acid  H2S206,  trithionic  acid  H2S306,  tetrathionic  acid 
H2S406,  and  pentathionic  acid  H2S6O6. 

SULPHURIC  ACID  H2S04 

Although  salts  of  sulphuric  acid,  such  as  calcium  sulphate  CaS04, 
are  exceedingly  plentiful  in  nature,  the  preparation  of  the  acid  by 
chemical  action  upon  the  salts  is  not  practicable.  The  sulphates, 
indeed,  interact  with  all  acids,  but  the  actions  are  reversible.  The 
completion  of  the  action  by  the  plan  used  in  making  hydrogen 


THE    OXIDES   AND    OXYGEN   ACIDS   OF   SULPHUR         281 

chloride  (p.  142),  involving  the  removal  of  the  sulphuric  acid  by 
distillation,  would  be  difficult  on  account  of  the  involatility  of  this 
acid.  It  boils  at  330°;  and  suitable  acids,  less  volatile  still,  which 
might  be  used  to  liberate  it,  do  not  exist.  We  are  therefore  com- 
pelled to  build  up  sulphuric  acid  from  its  elements. 

The  union  of  sulphur  dioxide  and  oxygen  by  the  contact  process, 
and  combination  of  the  trioxide  with  water  (p.  279),  is  the  best 
method  for  making  a  highly  concentrated  acid.  For  obtaining 
ordinary  "oil  of  vitriol/'  however,  the  "  chamber  process"  is  still 
used  extensively. 

Chemistry  of  the  Chamber  Process.  —  The  gases,  the  inter- 
actions of  which  result  in  the  formation  of  sulphuric  acid,  are: 
water  vapor,  sulphur  dioxide,  nitrous  anhydride  N2Os*  (see  index), 
and  oxygen.  These  are  obtained,  the  first  by  injection  of  steam, 
the  second  usually  by  the  burning  of  pyrite,  the  third  from  nitric 
acid  HN03,  and  the  fourth  by  the  introduction  of  air.  The  gases 
are  thoroughly  mixed  in  large  leaden  chambers,  and  the  sulphuric 
acid  forms  droplets  which  fall  to  the  floors.  In  spite  of  elaborate 
investigations,  instigated  by  the  extensive  scale  upon  which  the 
manufacture  is  carried  on  and  the  immense  financial  interests 
involved,  some  uncertainty  still  exists  in  regard  to  the  precise 
nature  of  the  chemical  changes  which  take  place.  According  to 
Lunge,  supporting  the  view  first  suggested  by  Berzelius,  the  greater 
part  of  the  product  is  formed  by  two  successive  actions,  the  first 
of  which  yields  a  complex  compound  that  is  decomposed  by  excess 
of  water  in  the  second: 

0-H 

H2O  +  2SO2  +  N2O3  +  O2  ->  2SO2  (  (1) 

X0  -  NO 

The  group  —  NO,  nitrosyl,  is  found  in  many  compounds.  Here,  if 
it  were  displaced  by  hydrogen,  sulphuric  acid  would  result.  Hence 
this  compound  is  called  nitrosylsulphuric  acid: 

0-H  /OH 

2SQ/  +H20^2S02^        +N203.  (2) 


X0-N0  OH 

*  This  gas  is  unstable,  breaking  up  in  part  into  nitric  oxide  NO  and  nitro- 
gen tetroxide  NO2:  N2Q3  <=*  NO  +  NO2.  In  this  process,  however,  the  mix- 
ture behaves  as  if  it  were  all  N2O3,  and  so  only  nitrous  anhydride  is  named  in 
this  connection. 


282  COLLEGE    CHEMISTRY 

The  equations  (1)  and  (2)  are  not  partial  equations  for  one  inter- 
action, but  represent  distinct  actions  which  can  be  carried  out 
separately.  In  a  properly  operating  plant,  indeed,  the  nitrosyl- 
sulphuric  acid  is  not  observed.  But  when  the  supply  of  water  is 
deficient,  white  "chamber  crystals,"  consisting  of  this  substance, 
collect  on  the  walls. 

The  explanation  of  the  success  of  this  seemingly  roundabout 
method  of  getting  sulphuric  acid  is  as  follows :  The  direct  union  of 
sulphur  dioxide  and  water  to  form  sulphurous  acid  is  rapid,  but  the 
action  of  free  oxygen  upon  the  latter,  2H2SO3  +  O2  — >  2H2S04,  is 
exceedingly  slow.  Reaching  sulphuric  acid  by  the  use  of  these  two 
changes,  although  they  constitute  a  direct  route  to  the  result,  is  not 
feasible  in  practice.  On  the  other  hand,  both  of  the  above  actions, 
(1)  and  (2),  happen  to  be  much  more  speedy,  and  so,  by  their  use, 
more  rapid  production  of  the  desired  substance  is  secured  at  the 
expense  of  a  slight  complexity. 

The  progress  of  the  first  action  is  marked  by  the  disappearance 
of  the  brown  nitrous  anhydride  and,  on  the  introduction  of  water, 
the  completion  of  the  second  stage  results  in  the  reproduction  of 
the  same  substance.  The  nitrous  anhydride  takes  part  a  large 
number  of  times  in  these  changes,  and  so  facilitates  the  conversion 
of  a  great  amount  of  sulphur  dioxide,  oxygen,  and  water  into  sul- 
phuric acid,  without  much  diminution  of  its  quantity.  Some  is 
lost,  however. 

The  loss  of  nitrous  anhydride  is  made  good  by  the  introduction 
of  nitric  acid  vapor  into  the  chamber.  This  acid  is  made  from  con- 
centrated sulphuric  acid  and  commercial  sodium  nitrate  NaNO3: 

jNaN03  +  H2S04  <=*  HN03T  +  NaHSO4. 

On  account  of  the  volatility  of  the  nitric  acid,  a  moderate  heat  is 
sufficient  to  remove  it  from  admixture  with  the  other  substances, 
and  its  vapor  is  swept  along  with  the  other  gases  into  the  apparatus. 
The  initial  action  which  the  nitric  acid  undergoes: 

H2O  +  2S02  +  2HN03  ->  2H2S04  +  N2O3, 
may  be  written,  to  show  the  anhydride  of  nitric  acid : 
H20  +  2SO2  +  H2O,N2O5  ->  2H2SO4  +  N2O3. 

The  two  molecules  of  water,  one  actually,  the  other  potentially, 
present,  with  the  two  molecules  of  sulphur  dioxide,  can  furnish  two 


THE   OXIDES  AND   OXYGEN  ACIDS   OF   SULPHUR         283 

molecules  of  sulphurous  acid  (H2S03).  The  N205  in  passing  to  the 
condition  N2O3  gives  up  the  two  units  of  oxygen  required  to  con- 
vert this  sulphurous  acid  into  sulphuric  acid. 

Details  of  the  Chamber  Process.  —  The  sulphur  dioxide  is 
produced  in  a  row  of  furnaces  A  (Fig.  87).  When  good  pyrite  is 
used,  the  ore  burns  unassisted  (p.  275),  while  impure  pyrite  and 
zinc-blende  ZnS  have  to  be  heated  artificially  to  maintain  the  com- 
bustion. The  gases  from  the  various  furnaces  pass  into  one  long 


FIG.  87. 

dust-flue,  in  which  they  are  mingled  with  the  proper  proportion  of 
air,  and  deposit  oxides  of  iron  and  of  arsenic,  and  other  materials 
which  they  transport  mechanically.  From  this  flue  they  enter  the 
Glover  tower  G,  in  which  they  acquire  the  oxides  of  nitrogen. 
Having  secured  all  the  necessary  constituents,  excepting  water,  the 
gases  next  enter  the  first  of  the  lead  chambers,  large  structures 
lined  completely  with  sheet  lead.  These  measure  as  much  as 
100  X  40  X  40  feet,  and  have  a  total  capacity  of  150,000  to  200,000 
cubic  feet.  As  the  gases  drift  through  these  chambers  they  are 
thoroughly  mixed,  and  an  amount  of  water  considerably  in  excess 
of  that  actually  required  is  injected  in  the  form  of  steam  at  various 
points.  The  acid,  along  with  the  excess  of  water,  condenses  and 


284  COLLEGE    CHEMISTRY 

collects  upon  the  floor  of  the  chamber,  while  the  unused  gases, 
chiefly  nitrous  anhydride  and  nitrogen,  the  latter  derived  from  the 
air  originally  admitted,  find  an  exit  into  the  Gay-Lussac  tower  L. 

This  is  a  tower  about  fifty  feet  in  height,  filled  with  tiles,  over 
which  concentrated  sulphuric  acid  continually  trickles.  The  object 
of  this  tower,  to  catch  the  nitrous  anhydride  and  enable  it  to  be 
reemployed  in  the  process,  is  accomplished  by  a  reversal  of  action 
(2)  above.  The  acid  which  accumulates  in  the  vessel  at  the  bottom 
of  this  tower  contains  the  nitrosylsulphuric  acid,  and  by  means  of 
compressed  air  is  forced  through  a  pipe  up  to  a  vessel  at  the  top  of 
the  Glover  tower  G.  When  this  "nitrous  vitriol"  is  mixed  with 
dilute  sulphuric  acid  from  a  neighboring  vessel,  by  allowing  both  to 
flow  down  into  the  tower,  the  nitrous  anhydride  is  once  more  set 
free  by  the  interaction  of  the  water  in  the  dilute  acid  (action  (2)). 
The  Glover  tower  is  filled  with  broken  flint  or  tiles,  and  the  heated 
gases  from  the  furnace  acquire  in  it  their  supply  of  nitrous  anhy- 
dride. Their  high  temperature  causes  a  considerable  concentra- 
tion of  the  diluted  sulphuric  acid  as  it  trickles  downward.  The 
acid,  after  traversing  this  tower,  is  sufficiently  strong  to  be  used 
once  more  for  the  absorption  of  nitrous  anhydride. 

To  replace  the  part  of  the  nitrous  anhydride  which  is  inevitably 
lost,  fresh  nitric  acid  is  furnished  by  small  open  vessels  N,  contain- 
ing sodium  nitrate  and  sulphuric  acid,  placed  in  the  flues  of  the 
pyrite-burners.  About  4  kg.  of  the  nitrate  are  consumed  for  every 
100  kg.  of  sulphur. 

The  acid  which  accumulates  upon  the  floors  contains  but  60  to 
70  per  cent  of  sulphuric  acid,  and  has  a  specific  gravity  of  1.5-1.62. 
The  excess  of  water  is  needed  to  facilitate  the  second  action.  It  is 
required  also  in  order  that  the  acid  upon  the  floor  may  not  after- 
wards absorb  and  retain  the  nitrous  anhydride,  for  this  substance 
combines  with  an  acid  containing  more  than  70  per  cent  of  hydro- 
gen sulphate. 

This  crude  sulphuric  acid  is  applicable  directly  in  some  chemical 
manufactures,  such  as  the  preparation  of  superphosphates  (q.v.). 
Concentration  is  effected  by  evaporation  in  pans  lined  with  lead, 
which  are  frequently  placed  over  the  pyrite-burners  in  order  to 
economize  fuel.  The  evaporation  in  lead  is  carried  on  until  a 
specific  gravity  1.7,  corresponding  to  77  per  cent  concentration,  is 
reached.  Up  to  this  point  the  sulphate  of  lead  formed  by  the 


THE    OXIDES   AND    OXYGEN   ACIDS   OF   SULPHUR         285 

action  of  the  sulphuric  acid  produces  a  crust  which  protects  the 
metal  from  further  action.  When  a  stronger  acid  is  required, 
the  water  is  driven  out  by  heating  the  sulphuric  acid  in  vessels  of 
glass  or  platinum,  or  even  of  cast  iron.  Iron  acts  upon  dilute 
sulphuric  acid,  displacing  the  hydrogen-ion,  but  not  upon  concen- 
trated sulphuric  acid,  which  is  not  ionized.  Commercial  sulphuric 
acid,  oil  of  vitriol,  has  a  specific  gravity  1.83-1.84,  and  contains 
about  93.5  per  cent  of  hydrogen  sulphate. 

Physical  Properties.  —  Pure  hydrogen  sulphate  has  a  sp.  gr. 
1.85  at  15°.  When  cooled,  it  crystallizes  (m.-p.  10.5°).  At  150°- 
180°  the  acid  begins  to  fume,  giving  off  sulphur  trioxide.  It 
boils  at  330°,  but  loses  more  sulphur  trioxide  than  water  and  finally 
yields  an  acid  of  constant  (p.  145)  boiling-point  (338°)  and  con- 
stant composition  (98.3.3  per  cent).  The  heat  of  solution  (p.  125) 
of  hydrogen  sulphate  is  very  great  (39,170  cal.).  The  solution  is 
thus  much  more  stable  (i.e.,  it  contains  much  less  energy)  than  the 
pure  substance,  and  hence  the  latter  absorbs  water  greedily. 

Commercial  sulphuric  acid  is  impure.  It  contains,  for  example, 
lead  sulphate,  which  appears  as  a  precipitate  when  the  acid  is 
diluted,  a,s  well  as  arsenic  trioxide  and  oxides  of  nitrogen  in  com- 
bination. 

Chemical  Properties  and  Uses  of  Hydrogen  Sulphate.— 

1.  The  compound  is  not  exceedingly  stable,  for  dissociation  into 
water  and  sulphur  trioxide  begins  far  below  the  boiling-point. 
The  vapor  of  the  acid  boiling  at  338°  contains  30  per  cent  of 
H20  +  80s,  which  recombine  when  the  vapor  is  condensed.  The 
dissociation  is  practically  complete  at  416°,  as  is  shown  by  the 
density  of  the  vapor.  When  raised  suddenly  to  a  red  heat  it  is 
broken  up  completely  into  water,  sulphur  dioxide,  and  oxygen. 

2.  When  sulphur  trioxide  is  dissolved  in  hydrogen  sulphate,  di- 
sulphuric  acid  H2S207,  a  solid  compound,  is  obtained.  Hydrogen 
sulphate  containing  80  per  cent  of  disulphuric  acid  is  known  as 
"  oleum,"  and  is  employed  in  chemical  industries.  The  salts  of 
disulphuric  acid  may  be  made  by  strongly  heating  the  acid  sul- 
phates, for  example: 

2NaHS04  <=*  Na2S2O7  +  H20|. 


286  COLLEGE    CHEMISTRY 

In  view  of  this  mode  of  preparation  by  the  aid  of  heat,  they  are 
frequently  known  as  pyrosulphates  (Gk.  irvp,  fire).  When  they  are 
dissolved  in  water,  the  acid  sulphates  are  reproduced. 

3.  With  salts  which  it  does  not  oxidize  (see  below),  hydrogen  sul- 
phate reacts  by  double  decomposition  and  sets  free  the  correspond- 
ing acid.    Where  the  new  acid  is  volatile,  as  in  the  case  of  hydrogen 
chloride  (p.  142),  we  are  furnished  with  one  of  the  cheapest  means 
of  preparing  acids.     Since  hydrogen  sulphate  is  dibasic  (p.  245), 
it  forms  both  acid  and  normal  salts,  such  as  NaHSC>4  and  Na2SO4. 
The  acid  sulphates  are  called  also  bisulphates,  because  they  con- 
tain twice  as  large  a  proportion  of  SO*  to  Na,  and  require  twice  as 
much  sulphuric  acid  for  their  preparation  as  do  the  neutral  sul- 
phates. 

4.  Sulphuric  acid  combines  vigorously  with 'water  to  form  at 
least  one  rather  stable  hydrate,  H2S04,H20  (m.-p.  8°).     On  this 
account,  sulphuric  acid  is  able  to  take  the  elements  of  water  from 
compounds  containing  hydrogen  and  oxygen,  especially  those  con- 
taining these  elements  in  the  proportion  2H  :  O.     Thus  paper, 
which  is  largely  cellulose  (CcHioOs)*,  wood  which  contains  much 
cellulose,  and  sugar  C^H^On  are  charred  by  it,  and  carbon  is 
set  free: 

<  CiaH^On  ->  12C  +  11H20. 

For  the  same  reason,  sulphuric  acid  is  used  in  drying  gases  with 
which  it  does  not  interact. 

5.  On  account  of  the  large  quantity  of  oxygen  which  hydrogen 
sulphate  contains,  and  its  instability  when  heated,  it  behaves  as 
an  oxidizing  agent.     This  property  has  already  been  illustrated  in 
connection  with  the  action  of  the  acid  upon  carbon,  sulphur,  and 
copper  (p.  276),  hydrogen  iodide  (p.  201),  and  hydrogen  bromide 
(p.  196).     The  sulphuric  acid  is  in  consequence  reduced  to  sulphur 
dioxide,  and  even  to  free  sulphur  or  hydrogen  sulphide.     The 
metals,  from  the  most  active  down  to  silver  (p.  260),  are  capable 
of  reducing  it,  the  sulphates*  being  formed.     The  more  active 
metals,  like  zinc,  reduce  it  to  hydrogen  sulphide  (p.  277),  the  less 

*  Note  that  the  sulphates,  and  not  the  oxides  of  the  metals  are  produced. 
Oxides  of  metals  could  not  be  formed  in  concentrated  sulphuric  acid,  because 
they  interact  with  the  latter  much  more  vigorously  than  do  the  metals,  to 
give  the  sulphates  (c/.  p.  146). 


THE    OXIDES   AND   OXYGEN   ACIDS   OF   SULPHUR         287 

Lctive,  like  copper,  give  sulphur  dioxide  (p.  276).     Hydrogen  is 
pot  liberated,  because  no  hydrogen-ion  is  present  in  concentrated 
(sulphuric  acid.     Gold  and  platinum  alone  do  not  interact  with  it. 
[Free  hydrogen  itself  is  oxidized  to  water  when  passed  into  hydrogen  ' 
[sulphate  at  160°:  S02(OH)2  +  H2  -»  SO2  +  2H20. 

Concentrated  sulphuric  acid  is  used  in  almost  all  chemical  in- 
dustries: for  example,  to  give  sodium  sulphate,  as  a  stage  in  the 
Le  Blanc  process  for  the  manufacture  of  soda;  in  the  refining  of 
I  petroleum;  in  the  manufacture  of  fertilizers,  such  as  superphos- 
phate; in  the  preparation  of  nitroglycerine  and  gun-cotton,  where 
it  assists  the  action  by  removing  water;  and  in  the  production  of 
coal-tar  dyes. 

Chemical  Properties  of  Aqueous  Hydrogen  Sulphate.  — 

The  solution  of  sulphuric  acid  H2S04,Aq  is  a  mixture,  whose  com- 
iponents  are:  undissociated  molecules  H2S04,  hydrogen-ion  H+, 
:hydrosulphate-ion  HS04~,  and  sulphate-ion  SO4=.  The  chemical 
properties  shown  by  the  solution  are  those  of  one  or  other  of  these 
components,  according  to  circumstances. 

Except  in  concentrated  solutions  (normal  or  stronger)  the  oxidiz- 
ing effects  of  the  undissociated,  molecular  substance  are  not 
encountered. 

The  presence  of  hydrogen-ion  is  shown  by  all  its  usual  properties 
(p.  246). 

Sulphate-ion  SO4=,  which  is  found  also  in  solutions  of  all  neutral 
and  acid  sulphates,  unites  with  all  positive  ions.  The  product, 
when  insoluble,  appears  as  a  precipitate.  The  introduction  of 
barium  ions,  for  example,  by  adding  a  solution  of  barium  nitrate 
or  chloride,  is  employed  as  a  test: 


Since  there  are  other  barium  salts  which  are  insoluble  in  water  (see 
Table  of  Solubilities),  but  no  common  ones  which  are  not  decom- 
posed by  acids,  dilute  nitric  acid  is  first  added  to  the  solution 
supposed  to  contain  the  sulphate-ion.  The  other  ions,  even  if 
present,  then  give  no  precipitate  with  barium-ion. 

Dilute  sulphuric  acid  is  used  for  many  purposes.  Thus,  it 
forms  the  liquid  in  the  lead  storage  battery,  and  is  employed  for 
cleaning  sheet  iron  before  tinning  and  galvanizing. 


288  COLLEGE    CHEMISTRY 

Sulphates,  —  The  acid  sulphates,  known  also  as  bisulphates, 
(see  p.  286),  may  be  produced  either  by  adding  to  dilute  sulphuric 
acid  half  an  equivalent  of  a  base,  and  evaporating:  NaOH  + 
H2S04 1=»  H20  +  NaHSO4,  or  by  actions  in  which  another  acid  is 
displaced,  as  in  making  hydrogen  chloride  (p.  141).  These  salts 
are  acid  in  reaction,  as  well  as  in  name  (cf.  p.  269),  because  HS04~, 
although  a  weak,  is  not  a  feeble  acid.  When  heated,  they  yield 
pyrosulphates  (p.  286). 

The  normal  (or  neutral)  sulphates  are  obtained  by  complete 
neutralization  and  evaporation,  or  by  the  second  of  the  above 
methods  when  a  sufficient  amount  of  the  salt  and  a  higher  tempera- 
ture are  used: 

NaCl  +  NaHSO4  <=±  Na^SO,  +  HClt- 

They  may  also  be  made  by  precipitation,  by -oxidation  of  a  sulphide 
at  a  high  temperature,  PbS  +  2O2  — >  PbSO4,  or  by  addition  of 
sulphur  trioxide  to  the  oxide  of  a  metal  (p.  280). 

Normal  sulphates  of  the  heavy  metals  decompose  at  a  red  heat, 
some  giving  off  sulphur  trioxide  (p.  279),  others  sulphur  dioxide 
and  oxygen.  The  sulphates  of  the  more  active  metals  and  of  lead, 
however,  are  not  affected  by  heating. 


OTHER  ACIDS  OF  SULPHUR 

Sulphurous  Acid  H2SO3,  Aq.  —  This  term  is  applied  to  the 
solution  of  sulphur  dioxide  in  water.  A  portion  of  the  sulphur 
dioxide  remains  dissolved  physically,  while  another  portion  is  in 
combination  with  the  water,  forming  sulphurous  acid.  This  in 
turn  is  ionized,  and  chiefly,  after  the  manner  of  the  weaker  dibasic 
acids,  into  two  ions,  H+  and  HSOs".  A  little  S03=  is  formed  from 
the  latter. 

Properties  of  Sulphurous  Acid.  —  The  acid  is  so  unstable  that 
it  cannot  be  obtained  excepting  in  solution  in  water.  Chemically  it 
is  a  comparatively  weak  acid.  As  a  reducing  agent,  it  is  slowly 
oxidized  to  sulphuric  acid  by  free  oxygen.  Sugar  and  glycerine 
act  as  negative  contact  agents  and  make  the  oxidation  much  slower. 
It  is  oxidized  more  rapidly  by  oxidizing  agents.  Thus,  when  free 


THE    OXIDES    AND    OXYGEN   ACIDS    OF   SULPHUR         289 

halogens  are  added  to  the  solution  (cf.  p.  161),  sulphuric  acid  and 
the  hydrogen  halide  are  formed: 

H2S03  +  HIO  «=>  H2S04  +  HI. 

Hydrogen  peroxide,  potassium  permanganate,  and  other  oxidizing 
agents  convert  the  substance  into  sulphuric  acid  likewise. 

Sulphurous  acid  has  the  power  of  uniting  directly  with  many 
organic  coloring  matters  and,  since  the  products  of  this  union  are 
usually  colorless,  it  is  employed  as  a  bleaching  agent.  It  is 
especially  useful  with  chemically  reactive  materials  like  silk,  wool, 
and  fragile  structures  like  straw,  which  are  likely  to  be  destroyed 
if  bleaching  powder  is  used.  The  compounds  thus  formed  are 
unstable,  and  lose  the  sulphurous  acid  again.  Hence,  wool  yellows 
with  age,  and  straw  hats  lose  their  whiteness.  As  a  disinfectant 
it  acts,  perhaps,  by  addition  likewise. 

As  a  dibasic  acid,  sulphurous  acid  forms  normal  salts  like 
Na2S03,  and  acid  salts  like  NaHS03. 

Consecutive  Reactions.  —  There  are  many  chemical  reactions 
that  proceed  in  two  stages,  which  can  be  carried  out  separately. 
This  is  the  case  with  the  two  reactions  used  in  the  chamber  process 
(p.  281).  The  actions  are  consecutive,  because  the  second  uses 
materials  produced  by  the  first.  It  may  be  noted  that  if  the 
second  action  is  as  speedy  as  the  first,  or  speedier,  then  no  inter- 
mediate products  will  be  detectable.  This  is  the  case  with  the 
chamber  process  reactions,  when  sufficient  steam  is  introduced, 
for  under  these  circumstances  no  solid  nitrosylsulphuric  acid  is 
deposited.  If  the  second  reaction  is  slower  than  the  first,  then  the 
products  of  the  first  reaction  will  accumulate,  and  become  notice- 
able. 

The  conception  of  consecutive  reactions  enables  us  to  under- 
stand and  remember  some  facts.  For  example,  it  was  mentioned 
that  when  dry  sulphur  is  oxidized,  we  obtain  sulphur  dioxide,  but 
when  moist  sulphur  is  oxidized,  by  the  air  or  otherwise,  the  only 
product  is  sulphuric  acid  (p.  267).  This  change  may  be  conceived 
of  as  proceeding  in  two  stages: 

S  +  O2  +  H20  -»  H2S03, 
2H2S03  +  02     ->2H2S04, 


290  COLLEGE    CHEMISTRY 

which  would  be  consecutive  reactions.  Since  oxidation  of  solid 
sulphur  can  proceed  only  on  the  surface,  it  is  slow.  Since  the 
sulphurous  acid  is  dissolved,  and  every  molecule  of  it  is  accessible 
to  the  dissolved  oxygen,  or  oxidizing  agent,  the  second  action 
should  be  speedier  and  consume  the  product  of  the  first  action  as 
fast  as  it  is  formed.  It  is,  therefore-,  quite  natural  that  no  sul- 
phurous acid  should  be  detectable  when  water  is  present. 

Sulphites.  —  The  acid  sulphites  of  the  alkali  metals,  KHS03 
and  NaHSOa,  when  in  solution,  are  acid  in  reaction,  owing  to  the 
appreciable  dissociation  of  the  ion  HSO3~.  The  sulphites  are 
readily  decomposed  by  acids  to  give  free  sulphurous  acid,  and  the 
latter  partly  decomposes,  yielding  sulphur  dioxide  (p.  275). 

Calcium  bisulphite  solution,  Ca(HSO3)2,  is  used  to  dissolve  the 
lignin  out  of  wood,  and  leave  the  pure  cellulose  (paper  pulp) 
employed  in  the  manufacture  of  paper  (q.v.). 

When  heated,  sulphites  undergo  decomposition.  The  sulphates, 
being  the  most  stable  of  all  the  salts  of  sulphur  acids,  are  formed 
when  the  salts  of  any  of  those  acids  are  decomposed  by  heating. 
The  nature  of  the  particular  salt  determines  what  other  products 
shall  appear.  Thus,  with  sodium  sulphite  Na^SOs,  one  molecule 
of  the  sulphite  furnishes  three  atoms  of  oxygen,  sufficient  to  oxi- 
dize three  other  molecules,  and  leaves  one  molecule  of  sodium 
sulphide  behind: 

3Na*S04. 


The  sulphites  are  as  readily  oxidized  as  is  the  acid  itself.  They 
are  slowly  converted,  both  in  solution  and  in  the  solid  form,  by  the 
influence  of  the  oxygen  of  the  air,  into  sulphates. 


Thiosulphuric  Acid  H^OQ.  —  This  acid  is  not  known  in  the 
free  condition,  but  its  salts  are  in  common  use  in  the  laboratory 
and  commercially.  Sodium  thiosulphate,  for  example,  is  pre- 
pared by  boiling  a  solution  of  sodium  sulphite  with  free  sulphur. 
The  action  is  something  like  the  addition  of  oxygen  to  sulphurous 
acid: 

NasSOs  +  S  -»  Na2S203    or    S03=  +  S  ->  S203=. 

Sodium  thiosulphate  ("hypo")  is  used  in  "fixing"  photographs. 


THE    OXIDES  AND   OXYGEN   ACIDS   OF   SULPHUR         291 

By  the  addition  of  acids  to  a  solution  of  sodium  thiosulphate, 
the  thiosulphuric  acid  is  set  free,  but  the  latter  instantly  decom- 
poses, giving  a  precipitate  of  sulphur: 

Na*S203  +  2HC1  *±  H2S203  +  2NaCl, 

H2S203  1=>  S|+  H2S03<=*  H20  +  S02T  . 


Persulphuric  Acid  -H^-SgOa.  —  This,  like  the  other  acids  just 
mentioned,  is  unstable,  and  can  be  kept  only  in  dilute  solution. 
Its  salts,  however,  are  coming  into  use  for  commercial  purposes 
and  for  "reducing"  negatives  in  photography.  The  salts  are 
prepared  by  electrolyzing  sodium-hydrogen  sulphate  NaHSCX  in 
concentrated  solution  (Hugh  Marshall).  The  persulphuric  acid, 
formed  by  the  union  of  the  negative  ions  in  pairs  as  they  are 
discharged  on  the  anode, 


undergoes  double  decomposition  with  the  excess  of  sodium  bisul- 
phate,  and  the  less  soluble  sodium  persulphate  crystallizes  out. 
The  other  salts  are  made  by  double  decomposition  from  this  one. 

Compounds  of  Sulphur  and  Chlorine.  —  When  chlorine.  gas 
is  passed  over  heated  sulphur,  it  is  absorbed  and  sulphur  mono- 
chloride,  a  reddish-yellow  liquid,  boiling  at  138°,  is  obtained.  The 
molecular  weight  of  this  substance,  as  shown  by  the  density  of  its 
vapor,  indicates  that  it  possesses  the  formula  S2C12.  When  thrown 
into  water,  it  is  rapidly  hydrolyzed,  producing  sulphur  dioxide  and 
sulphur:  2S2C12  +  2H20  ->  S02  +  4HC1  +  3S. 

Sulphur  itself  dissolves  very  freely  in  the  monochloride,  and  the 
solution  is  employed  in  vulcanizing  rubber. 

Sulphur  dioxide  and  chlorine  gases,  when  exposed  to  direct  sun- 
light, unite  to  form  a  liquid  known  as  sulphuryl  chloride  SO2C12. 
Camphor  causes  the  union  to  take  place  much  more  rapidly,  owing 
to  some  catalytic  effect.  The  compound  is  a  colorless  liquid,  boil- 
ing at  69°.  With  water  it  gives  sulphuric  acid  and  hydrogen 
chloride  : 

S02C12  +  2H2O  -*  SO2(OH)2  +  2HC1. 

Graphic  Formula  of  Sulphuric  Acid.  —  The  actions  just 
mentioned  give  a  clue  to  the  constitution  of  sulphuric  acid.    Since 
*  See  footnote,  p.  276. 


292  COLLEGE   CHEMISTRY 

chlorine  does  not  combine  directly  with  oxygen,  but  does  com- 
bine readily  with  sulphur,  we  may  assume  that,  in  the  formation 
of  sulphuryl  chloride,  S02  +  C12  —  >  SO2C12,  the  chlorine  unites 
more  intimately  with  the  sulphur  in  the  molecule  S02  : 


The  action  of  water  upon  the  product  is  presumably  similar  to 
that  of  water  on  phosphorus  tribromide  (p.  197): 

(K     !'"ci  .......  H'4-O-H  0.        0-H 


cr    >ci      Hfo-H  cr    NO-H 

The  last  is  called  the  structural  formula  of  sulphuric  acid.  It  is  not 
thereby  implied  that  the  atoms  in  its  molecules  are  attached  pre- 
cisely in  this  manner,  however,  but  rather  that  the  chemical  be- 
rhavior  of  the  substance,  as  being  partly  an  oxide  and  partly  an 
hydroxide  of  sulphur,  is  symbolized  in  this  fashion.  Such  graphic 
formulae  are  of  great  value  in  expressing  the  chemical  behavior  of 
the  complex  compounds  of  carbon. 

Exercises.  —  1.   What  ground  is  there  for  assigning  the  formula 
instead  of  S204  to  sulphur  dioxide  (p.  277)? 

2.  Explain  why  nitric  acid  is  completely  displaced  by  the  action 
of  sulphuric  acid  on  sodium  nitrate  (p.  282). 

3.  What  are  the  relative  volumes,  (a)  of  sulphur  dioxide  and 
nitrogen  (p.  150)  resulting  from  the  roasting  of  pyrite  (p.  275),  (6) 
of  air  and  sulphur  dioxide  in  making  sulphuric  acid,  (a)  of  nitrogen 

left)  to  sulphur  dioxide  (used)  in  making  sulphurid  acid,  when 
yriteis  the  souroo£ • — — 


^  4.   Make  a  list  of,  and  classify,  the  various  applications  of 
sulphuric  acid  to  the  liberation  of  other  acids. 

5.  Formulate  the  behavior  of  the  hydrosulphate-ion  (p.  287) 
phen  a  solution  of  barium  chloride  is  added  to  a  rather  concen- 
trated solution  of  sulphuric  acid. 

6.  Assign  to  the  proper  class  of  ionic  actions  (pp.  259,  270),  (a) 
the  action  of  iodine  on  sulphurous  acid  (p.  289),  (6)  of  sulphur  on 
sodium  sulphite  (p.  290),  (c)  the  formation  of  persulphuric  acid 


*«+** 


CHAPTER  XXII 

SELENIUM   AND   TELLURIUM 
THE  CLASSIFICATION  OF  THE  ELEMENTS 

ALONG  with  sulphur,  chemists  group  two  other  elements,  sele- 
nium (Se,  at.  wt.  79.2)  and  tellurium  (Te,  at.  wt.  127.5).  If,  while 
this  and  the  next  page  are  read,  the  nature  of  the  chief  compounds 
of  sulphur  is  kept  in  mind,  the  analogy  between  the  nature  and 
chemical  behavior  of  the  three  elements  and  their  corresponding 
compounds  will  be  obvious  (see  Chemical  relations  of  the  sulphur 
family,  below). 

Occurrence  and  Properties  of  Selenium  Se.  —  Selenium 
(Gk.,  the  moon)  occurs  free  in  some  specimens  of  native  sulphur, 
and  in  combination  often  takes  the  place  of  a  small  part  of  the 
sulphur  in  pyrite  (FeS2).  It  is  found  free  in  the  dust-flues  of  the 
pyrite-burners  of  sulphuric  acid  works.  The  familiar  forms  are, 
the  red  precipitated  variety,  which  is  amorphous  and  soluble  in 
carbon  disulphide,  and  the  lead-gray, .  semi-metallic  variety,  ob- 
tained by  slow  cooling  of  melted  selenium,  which  is  insoluble,  and 
melts  at  217°.  In  the  latter  form  it  has  some  capacity  for  con- 
ducting electricity,  which  is  greatly  increased  by  exposure  to  light 
in  proportion  to  the  intensity  of  the  illumination.  A  photometer, 
using  this  property,  has  been  devised  by  Joel  Stebbins  (1914),  for 
measuring  the  relative  intensity  of  the  light  of  different  stars. 
Selenium  boils  at  680°,  and  at  high  temperatures  has  a  vapor 
density  corresponding  to  the  formula  Sea. 

The  element  combines  directly  with  many  metals,  burns  in 
oxygen  to  form  selenium  dioxide,  and  unites  vigorously  with 
chlorine. 

Compounds  of  Selenium.  —  Ferrous  selenide,  made  by  heat- 
ing iron  filings  with  selenium  (cf.  p.  13),  when  treated  with  con- 
centrated hydrochloric  acid  gives  hydrogen  selenide: 

FeSe  +  2HC1  fc»  H2Se  t  +  FeCl«. 
293 


294  COLLEGE    CHEMISTRY 

The  compound  is  a  poisonous  gas,  which  possesses  an  odor  recalling 
rotten  horse-radish,  and  is  soluble  in  water.  The  solution  is  faintly 
acid  in  reaction,  and  deposits  seleniiun  when  exposed  to  the  action 
of  the  air  (cf.  p.  270).  Other  selenides,  which,  with  the  exception 
of  those  of  potassium  and  sodium,  are  insoluble  in  water,  may  be 
precipitated  by  leading  the  gas  into  solutions  of  soluble  salts  of 
appropriate  metals  (cf.  p.  273). 

The  dioxide  Se02  is  a  solid  body  formed  by  burning  selenium. 
Selenious  acid  H2SeO3  may  be  made  by  dissolving  the  dioxide  in  hot 
water,  or  by  oxidizing  selenium  with  boiling  nitric  acid.  Unlike 
sulphur  (p.  267),  the  element  gives  little  of  the  higher  acid  H2Se04 
by  this  treatment.  The  acid  is  reduced  by  sulphurous  acid  to 
selenium:  H2Se03  +  2H2S03  ->  2H2S04  +  H20  +  Se. 

No  trioxide  is  known.  Selenic  acid  H2Se04,  a  white  solid,  is 
made  in  solution  by  oxidizing  silver  selenite  with  bromine-water 
(which  contains  hypobromous  acid,  cf.  p.  161),  and  filtering: 


+  HBrO  ->  Ag2Se04  +  HBr. 
2HBr  +  Ag2Se04  -»  2AgBr  |  +  H2Se04 
Br2  +  H20  +  Ag2SeO3  ->  2AgBr  J,  +  H2Se04.' 


It  is  itself  a  powerful  oxidizing  agent  and,  even  in  dilute  solution. 
liberates  chlorine  from  hydrochloric  acid:  H2Se04  +  2HC1  —  > 
H2Se03  +  H20  +  C12.  Sulphuric  acid  (cf.  p.  286),  on  the  other 
hand,  is  an  oxidizing  agent  only  in  somewhat  concentrated  form, 
and  even  then  it  can  oxidize  hydrobromic  acid  (p.  196),  but  not 
hydrochloric  acid. 

Tellurium  Te.  —  Tellurium  (Lat.,  the  earth)  occurs  in  sylvan- 
ite  in  combination  with  gold  and  silver.  It  is  a  white,  metallic, 
crystalline  substance,  melting  at  452°  (b.-p.  1400°).  The  free 
element  unites  with  metals  directly,  and  burns  in  air  to  form  the 
dioxide. 

The  compounds  of  tellurium  are  similar  in  composition  and  mode 
of  preparation  to  those  of  selenium.  Some  differences  in  chemical 
behavior  are  significant,  however.  Thus,  tellurious  acid  H2TeO3  is 
a  very  feeble  acid  and  is  also  somewhat  basic,  a  sulphate  (2Te02, 
S03)  and  a  nitrate  (Te203(OH)N03)  being  known.  In  this  respect 


THE    PERIODIC    SYSTEM  295 

it  differs  markedly  from  sulphurous  acid.  Telluric  acid  does  not 
affect  indicators,  and  is  therefore  actually  more  feebly  acidic  than 
is  hydrogen  sulphide.  Tellurium  tetrachloride  TeClj,  although 
hydrolyzed  by  water,  exists  in  solution  with  excess  of  hydrogen 
chloride:  TeCU  +  3H20  <=>  H2Te03  +  4HC1,  showing  the  telluri- 
ous  acid  to  be  basic  in  properties  and  the  element  tellurium  to  be, 
to  a  certain  degree,  a  metallic  element. 

The  Chemical  Relations  of  the  Sulphur  Family.  —  It  will 
be  seen  that  sulphur,  selenium,  and  tellurium  are  bivalent  elements 
when  combined  with  hydrogen  or  metals.  In  combination  with 
oxygen  they  form  unsaturated  compounds  of  the  form  XIV02,  while 
their  highest  valence  is  found  in  SO3,  TeOs,  and  H2SeO4,  where 
they  must  be  sexivalent.  The  general  behavior  of  corresponding 
compounds  is  very  similar.  At  the  same  time,  there  is  in  all  cases 
a  progressive  change  as  we  proceed  from  sulphur  through  selenium 
to  tellurium.  The  elementary  substances  themselves,  for  example, 
become  more  like  metals,  physically,  and  they  show  higher  and 
higher  melting-points.  The  affinity  for  hydrogen  decreases,  as  is 
shown  by  the  increasing  ease  with  which  the  compounds  H2X  are 
oxidized  in  air.  The  affinity  for  oxygen  likewise  decreases,  for  the 
elements  become  increasingly  difficult  to  raise  to  the  highest  state 
of  oxidation.  On  the  other  hand,  the  tendency  to  form  higher 
chlorides  becomes  greater.  We  note  also  that  the  compounds 
H2X04  become  less  and  less  active  as  acids,  and  that  a  basic 
tendency  begins  to  assert  itself. 

THE  PERIODIC  SYSTEM 

Classification,  or  the  arrangement  of  facts  on  the  basis  of  like- 
ness, is  part  of  the  method  of  science.  It  is  needed  to  make 
possible  the  systematic  description  of  the  ascertained  facts,  which 
is  a  great  aid  to  the  memory,  and  to  furnish  a  guide  in  investigation, 
by  suggesting  relations  and  so  pointing  out  directions  in  which  new 
facts  of  interest  may  be  found.  Thus,  as  an  aid  to  memory,  we 
have  treated  the  halogens  as  one  family  and  S,  Se,  and  Te  as  an- 
other. In  each  case,  we  have  presented  the  properties  common  to 
all  members  of  the  group,  and  have  then  pointed  out  the  differ- 
ences. Again,  in  investigation,  as  soon  as  we  have  discovered  that 


296  COLLEGE    CHEMISTRY 

sulphur  and  selenium  are  allied  elements,  we  realize  the  direction 
in  which  fruitful  results  may  be  expected,  and  we  proceed  to  make 
the  corresponding  compounds  and  to  note  the  resemblances  and 
differences  in  the  conditions  for  preparation  and  in  the  properties 
of  the  compounds  obtained. 

Metallic  and  Non-Metallic  Elements.  —  Thus  far  we  have 
found  the  division  into  metallic  and  non-metallic  elements  very 
serviceable  for  classification  in  terms  of  chemical  relations  (p.  163). 
This  distinction  we  shall  continue  to  employ.  The  metallic,  or 
positive  elements  (p.  94),  (1)  form  positive  radicals  and  ions  con- 
taining no  other  element  (c/.  p.  247).  Thus  the  metals  give  sul- 
phates, nitrates,  carbonates,  and  other  salts,  which  furnish  a 
metallic  ion,  such  as  Na+  or  K+,  together  with  the  ions  804=,  N03~ 
and  C03=.  (2)  Their  hydroxides,  KOH,  Ca(OH)2,  etc.,  give  the 
same  metallic  ion,  and  the  rest  of  the  molecule  forms  hydroxide-ion. 
That  is  to  say,  their  hydroxides  are  bases  and  their  oxides  are 
basic.  The  metallic  elements  often  enter,  but  only  with  other 
elements,  into  the  composition  of  a  negative  ion,  as  is  the  case  with 
manganese  in  K.Mn04,  with  chromium  in  K2.Cr207,  and  with  silver 
in  KAg(CN)2. 

The  non-metallic  or  negative  elements  (1)  are  found  chiefly  in 
negative  radicals  and  ions.  They  form  no  nitrates,  sulphates,  car- 
bonates, etc.,  for  they  could  not  do  so  without  themselves  alone 
constituting  the  positive  ion.  We  have  no  such  salts  of  sulphur, 
carbon,  or  phosphorus,  for  example.  (2)  Their  hydroxides,  al- 
though their  formulae  may  be  written  C1O2OH,  P(OH)3,  SO2(OH)2, 
furnish  no  hydroxyl  ions,  as  this  would  involve  the  same  conse- 
quence. These  hydroxides  are  divided  by  dissociation,  in  fact,  so 
that  the  non-metal  forms  part  of  a  compound  negative  radical,  and 
the  other  ion  is  hydrogen-ion,  C103.H,  P03H.H2,  S04.H2.  Their 
oxides  are  acidic.  (3)  Their  halogen  compounds,  like  PBr3  (p.  197) 
and  S2C12  (p.  291),  are  completely  hydrolyzed  by  water,  and  the 
actions  are  not,  in  general,  reversible.  The  halides  of  the  typical 
metals  are  not  hydrolyzed  (see  Chap.  XXXIII),  and  with  those  that 
are  not  typical,  the  action  is  reversible. 

The- distinction  is  not  perfectly  sharp,  however.  Thus,  zinc 
gives  both  salts  like  the  sulphate,  Zn.S04,  and  chloride,  Zn.Cl2,  and 
compounds  like  sodium  zincate  (p.  56),  ZnO2.Na2. 


THE    PERIODIC    SYSTEM  297 

Classification  by  Atomic  Weights.  —  Newlands  (1863-4)  dis- 
covered a  surprising  regularity  that  became  apparent  when  the  ele- 
ments were  placed  in  the  order  of  ascending  atomic  weight. 
Omitting  hydrogen  (at.  wt.  1)  the  first  seven  were:  lithium  (7), 
glucinum  (9),  boron  (11),  carbon  (12),  nitrogen  (14),  oxygen  (16), 
fluorine  (19).  These  are  all  of  totally  different  classes,  and  include 
first  a  metal  forming  a  strongly  basic  hydroxide,  then  a  metallic 
element  of  the  less  active  sort,  then  five  non-metals  of  increasingly 
negative  character,  the  last  being  the  most  active  non-metal 
known.  The  next  element  after  fluorine  (19)  was  sodium  (23), 
which  brings  us  back  sharply  to  the  elements  that  form  strongly 
basic  hydroxides.  Omitting  none,  the  next  seven  elements  were 
sodium  (23),  magnesium  (24.4),  aluminium  (27),  silicon  (28.4), 
phosphorus  (31),  sulphur  (32),  chlorine  (35.5).  In  this  series  there 
are  three  metals  of  diminishing  positiveness,  followed  by  four  non- 
metals  of  increasing  negative  activity,  the  last  being  a  halogen  very 
like  fluorine.  On  account  of  the  fact  that  each  element  resembles 
most  closely  the  eighth  element  beyond  or  before  it  in  the  list,  the 
relation  was  called  the  law  of  octaves.  After  chlorine  the  octaves 
become  less  easy  to  trace. 

That  this  periodicity  in  chemical  nature  is  more  than  a  coinci- 
dence is  shown  by  the  fact  that  the  valence  and  even  the  physical 
properties,  such  as  the  specific  gravity,  show  a  similar  fluctuation 
in  each  series.  In  the  first  two  series  the  compounds  with  other 
elements  are  of  the  types: 

LiCl,   G1C12,   BO,  CO,  '  0H2,  FH.          ; 


Thus  the  valence  towards  chlorine  or  hydrogen  ascends  to  four  and 
then  reverts  to  one  in  each  octave.  The  highest  valence,  shown  in 
oxygen  compounds,  ascends  from  lithium  to  nitrogen  with  values 
one  to  five,  and  then  fails  because  compounds  are  lacking.  In  the 
second  octave,  however,  it  goes  up  continuously  from  one  to 
seven. 

Again,  the  specific  gravities  of  the  elements  in  the  second  series, 
using  the  data  for  red  phosphorus  and  liquid  chlorine,  are: 

Na  0.97,  Mg  1.75,  Al  2.67,  Si  2.49,  P  2.14,  S  2.06,  Cl  1.33. 


298  COLLEGE    CHEMISTRY 

Mendelejeff's  Scheme.  —  In  1869  Mendelejeff  published  an 
important  contribution  towards  adjusting  the  difficulty  which  the 
elements  following  chlorine  presented,  and  developed  the  whole 
conception  so  completely  that  the  resulting  system  of  classification 
has  been  connected  with  his  name  ever  since.  Almost  simultane- 
ously Lothar  Meyer  made  similar  suggestions,  but  did  not  urge 
them  with  the  same  conviction  or  elaborate  them  so  fully.  The 
table  on  the  following  page,  in  which  the  atomic  weights  are  ex- 
pressed in  round  numbers,  is  a  modification  of  one  of  Mendelejeff's. 

The  chief  change  from  the  arrangement  in  simple  octaves  is  that 
the  third  series,  beginning  with  potassium,  is  made  to  furnish 
material  for  two  octaves,  potassium  to  manganese  and  copper  to 
bromine,  and  is  called  a  long  series.  The  valences  fall  in  with  this 
plan  fairly  well.  Copper,  while  usually  bivalent,  forms  also  a 
series  of  compounds  in  which  it  is  univalent.  Iron,  cobalt,  and 
nickel  cannot  be  accommodated  in  either  octave,  as  their  valences 
are  always  two  or  three.  At  the  time  Mendelejeff  made  the  table, 
three  places  in  the  third  long  series  had  to  be  left  blank,  as  a  tri- 
valent  element  [Sc]  was  lacking  in  the  first  octave  of  the  series,  and 
a  trivalent  [Ga]  and  a  quadrivalent  one  [Ge]  in  the  second.  These 
places  have  since  been  filled,  as  we  shall  presently  see.  The  first 
two  (the  short)  series  have  been  split  in  the  table,  as  lithium  and 
sodium  closely  resemble  potassium,  while  the  remaining  members 
of  these  series  fall  more  naturally  over  the  corresponding  elements 
of  the  second  octave  of  the  third  series. 

The  fourth  series  (long)  is  nearly  complete.  It  begins  with  an 
active  alkali  metal,  rubidium,  and  ends  with  iodine,  a  halogen. 
The  rule  of  valence  is  strictly  preserved  throughout  the  series,  and 
in  general  the  elements  fall  below  those  which  they  most  closely 
resemble. 

The  fifth,  sixth,  and  seventh  (long)  series  are  incomplete,  but  the 
order  of  the  atomic  weights  and  the  valence  enable  us  satisfac- 
torily to  place  those  elements  which  are  known.  The  chemical 
relations  to  elements  of  the  fourth  series  justify  the  position  as- 
signed to  each.  Caesium,  for  example,  is  the  most  active  of  the 
alkali  metals;  barium  has  always  been  classed  with  strontium,  and 
bismuth  with  antimony. 

In  two  cases  a  slight  displacement  of  the  order  according  to 
atomic  Weights  is  necessary.  Cobalt  is  put  before  nickel  because  it 


THE    PERIODIC    SYSTEM 


299 


Oi 

~£ 


OS 


02 


O05 


2 


Is 


OS 


0)0 

OS 


o3o 

0° 


300  COLLEGE   CHEMISTKY 

resembles  iron  more  closely.  Tellurium  and  iodine  are  placed  in 
that  order  to  bring  them  into  the  sulphur  and  halogen  groups 
respectively.  Their  valence  and  other  chemical  relations  both 
require  this.  The  general  agreement,  however,  is  very  remarkable. 

General  Relations  in  the  System.  —  In  every  octave  the 
valence  towards  oxygen  ascends  from  one  to  seven,  while  that 
towards  hydrogen,  in  the  cases  of  the  last  four  elements  (when  they 
combine  with  hydrogen  at  all),  descends  from  four  to  one.  The 
physical  properties  fluctuate  within  the  limits  of  each  series  in  a 
similar  way.  The  values  of  each  physical  constant  for  correspond- 
ing members  of  the  successive  series  do  not  exactly  coincide,  how- 
ever. A  progressive  change,  as  we  descend  each  vertical  column, 
is  the  rule.  Thus  the  specific  gravities  (water  =  1)  of  the  alkali 
metals  rise  from  lithium  (0.53)  to  caesium  (1.87).  In  the  same 
group  the  melting-points  descend  from  lithium  (186°)  to  caesium 
(26.5°). 

As  yet  no  exact  mathematical  (quantitative)  relation  between 
the  values  for  any  property  and  the  values  of  the  atomic  weights 
has  been  discovered;  only  a  general  (qualitative)  relationship  can 
be  traced.  Anticipating  the  discovery  of  some  more  exact  mode 
of  stating  the  relationship  in  each  case,  and  remembering  that 
similar  values  of  each  property  recur  periodically,  usually  at  inter- 
vals corresponding  to  the  length  of  an  octave  or  series,  the  principle 
which  is  assumed  to  underlie  the  whole,  the  periodic  law,  is  stated 
thus :  All  the  properties  of  the  elements  are  periodic  functions  of 
their  .atomic  weights. 

That  the  chemical  relations  of  the  elements  vary  just  as  do  the 
physical  properties  of  the  simple  substances  is  easily  shown. 
Thus,  each  series  begins  with  an  active  metallic  (positive)  element, 
and  ends  with  an  active  non-metallic  (negative)  element,  the  inter- 
vening elements  showing  a  more  or  less  continuous  variation 
between  these  limits.  Again,  the  elements  at  the  top  are  the  least 
metallic  of  their  respective  columns.  As  we  descend,  the  members 
of  each  group  are  more  markedly  metallic  (in  the  first  columns),  or, 
what  is  the  same  thing,  less  markedly  non-metallic  (in  the  later 
columns;  cf.  p.  296). 

In  the  first  series  boron  is  the  first  non-metal  we  encounter.  In 
the  second  series  silicon  is  the  first  such  element.  In  the  third 


THE   PERIODIC   SYSTEM  301 

there  is  more  difficulty  in  deciding.  Titanium,  vanadium,  and 
germanium  are  usually,  though  with  questionable  propriety,  classed 
as  metallic  elements.*  Selenium  is  undoubtedly  a  non-metal. 
Arsenic  is,  on  the  whole,  a  non-metal.  In  the  fourth  series  telu- 
rium  is  commonly  considered  to  be  the  first  non-metal.  Thus  a 
zigzag  line,  shown  in  the  table,  separates  all  the  non-metals  from 
the  rest  of  the  elements,  and  confines  them  in  the  right-hand  upper 
corner. 

A  more  compact  form  of  the  table  is  printed  at  the  end  of  this 
book,  opposite  the  rear  cover.  The  only  difference  between  this 
and  the  other  is  that  the  two  octaves  of  each  long  series  have  been 
placed  in  the  same  set  of  seven  main  columns.  The  iron,  palla- 
dium, and  platinum  groups  occupy  a  column  on  the  right  of  the 
main  columns,  and  are  often  called  collectively  the  eighth  group. 
The  newly  discovered  elements,  found  chiefly  in  the  air,  have  been 
placed  at  the  left-hand  side.  Since  they  do  not  enter  into  com- 
bination at  all,  their  valence  may  appropriately  be  given  as  zero. 
With  the  exception  of  argon,  the  values  of  their  atomic  weights 
agree  well  with  this  assignment.  Hydrogen  is  the  only  common 
element  whose  place  is  still  in  debate.  The  valence  is  shown  by 
the  general  formulae  at  the  head  of  each  column. 

Applications  of  the  Periodic  System.  —  The  system  has 
found  application  chiefly  in  four  ways: 

1.  In  the  prediction  of  new  elements.  Mendelejeff  (1871)  drew 
attention  to  the  blank  then  existing  between  calcium  (40)  and  tita- 
nium (48).  He  predicted  that  an  element  to  fit  this  place  would 
have  an  atomic  weight  44  and  would  be  trivalent.  From  the 
nature  of  the  surrounding  elements,  he  very  cleverly  deduced  many 
of  the  physical  and  chemical  properties  of  the  unknown  element  and 
of  its  compounds.  In  1879  Nilson  discovered  scandium  (44),  and 
its  behavior  corresponded  closely  with  that  predicted.  Mendele- 
jeff described  accurately  two  other  elements,  likewise  unknown  at 
the  time.  In  1875  Lecoque  de  Boisbaudran  found  gallium,  and 
in  1888  Winkler  discovered  germanium,  and  these  blanks  were 
filled. 

*  In  discussing  chemical  relations,  the  term  metallic  element  is  preferable 
to  metal.  The  free  element  (e.g.,  arsenic)  may  have  the  luster  of  a  metal, 
and  yet  the  element,  in  combination,  may  be  non-metallic  or  negative. 


302  COLLEGE    CHEMISTRY 

2.  By  enabling  us  to  decide  on  the  correct  values  for  the  atomic 
weights  of  some  elements,  when  the  equivalent  weights  have  been 
measured,  but  no  volatile  compound  is  known  (cf.  pp.  104  and  118). 
Thus,  the  equivalent  weight  of  indium  was  38  and,  as  the  element 
was  supposed  to  be  bivalent,  it  received  the  atomic  weight  76.     It 
was  quite  out  of  place  near  arsenic  (75),  however,  being  decidedly 
a  metal.    As  a  trivalent  element  with  the  atomic  weight  115,  it  fell 
between  cadmium  and  tin.     Later  work  fully  justified  the  change. 
Quite  recently,  radium  has  been  discovered,  and  found  to  have 
the  equivalent  weight   113  and  to  resemble  barium.     If,   like 
barium,  it  is  bivalent,  it  occupies  a  place  under  this  element,  in  the 
last  series. 

3.  By  suggesting  problems  for  investigation.    The  periodic  system 
has  been  of  constant  service  in  the  course  of  inorganic  research,  and 
has  often  furnished  the  original  stimulus  to  such  work  as  well.     For 
example,  the  atomic  weight  of  tellurium  bore  the  value  128  when 
the  table  was  first  constructed,  and  it  was  confidently  expected  that 
reexamination  would  bring  this  value  below  that  of  iodine  (then 
127,  now  126.92).     Several  most  careful  studies  of  the  subject  have 
been  made  by  different  methods.     It  seems  probable  that  the  real 
value  of  the  atomic  weight  is  not  far  from  Te  =  127.5,  and  there- 
fore more  than  half  a  unit  greater  than  that  of  iodine.     Since,  how- 
ever, mathematical  correspondence  is  found  nowhere  in  the  system, 
the  existence  of  marked  inconsistencies  like  this  need  not  shake  our 
confidence  in  its  value  when  it  is  used  with  due  consideration  of  the 
degree  of  correspondence  to  be  expected. 

In  the  same  way,  incorrect  values  of  many  physical  properties 
have  been  detected,  and  have  been  rectified  by  more  careful 
work. 

4.  By  furnishing  a  comprehensive  classification  of  the  elements, 
arranging  them  so  as  to  exhibit  the  relationships  among  the  physical 
and  chemical  properties  of  the  elements  themselves  and  of  their 
compounds.     Constant  use  will  be  made  of  this  property  of  the 
table  in  the  succeeding  chapters.     Having  disposed  of  the  halogen 
and  sulphur  families  (excepting  the  oxygen  compounds  of  the 
former),  situated,  respectively,  in  the  seventh  and  sixth  columns 
of  the  table  (at  the  end  of  this  book),  we  shall  presently  take  up 
nitrogen  and  phosphorus  from  the  right  side  of  the  fifth  column. 
Then  from  the  fourth  column,  we  shall  select  carbon  and  silicon, 


THE    PERIODIC    SYSTEM  303 

and  from  the  third  boron,  leaving  the  other,  more  decidedly  metal- 
lic elements  for  later  treatment. 

Moseley's  Atomic  Numbers.  —  We  have  seen  that  simple, 
mathematical  relations  between  the  atomic  weights  and  the  physi- 
cal or  chemical  properties  of  an  element  do  not  exist.  In  several 
instances,  the  atomic  weights  are  not  even  in  the  same  order  as 
are  the  values  of  the  properties.  We  have  now  obtained  from 
another  direction  numbers  which  seem  to  be  more  fundamental 
even  than  atomic  weights. 

Visible  light,  X-rays,  and  wireless  electric  waves  are  all  vibra- 
tions of  the  same  nature  in  the  ether.  They  differ  only  in  wave- 
length, the  order  of  the  wave-lengths  being  10~5  cm.,  10~8  cm.,  and 
106  cm.  (10  kilometers),  respectively.  Now,  just  as  the  spectrum 
of  visible  light  is  obtained  by  using  a  grating,  on  which  the  rulings 
are  separated  by  distances  of  the  order  of  the  wave-length  of  such 
light,  so  ordinary  crystals  give  spectra  of  X-rays,  because  they  are 
composed  of  particles  arranged  in  rows  about  one  thousand  times 
closer  and  so  form  a  suitable  grating  for  X-rays.  This  fact  was 
first  discovered  by  Dr.  Laue  of  the  University  of  Zurich  (1912). 
The  X-rays  are  produced  in 
an  evacuated  tube  by  cathode 
rays,  which  are  streams  of 
electrons  emanating  from 
the  cathode  (C,  Fig.  88), 
when  they  strike  the  anti-  FlG  88 

cathode  (A). 

With  different  elements  on  the  anti-cathode,  X-rays  of  slightly 
different  wave-lengths,  and  therefore  giving  different  X-ray 
spectra,  are  produced.  By  using  different  elements,  Moseley 
(1914)  has  found  that  the  higher  the  atomic  weight  the  shorter  the 
wave-length  of  the  characteristic  X-rays.  When  the  elements 
are  arranged  in  the  order  of  these  wave-lengths,  whole  numbers 
can  be  assigned  to  each  which  are  inversely  proportional  to  the 
wave-lengths  of  corresponding  lines  in  their  X-ray  spectra.  These 
atomic  numbers  have  been  determined  for  most  of  the  elements,  the 
atomic  weights  of  which  lie  between  those  of  aluminium  and  gold. 
In  the  following  table,  the  atomic  numbers  for  these  elements  are 
given  and,  for  the  sake  of  greater  completeness,  numbers  for  the 
twelve  elements  preceding  Al  have  been  inserted  also. 


304 


COLLEGE    CHEMISTRY 


ATOMIC  NUMBERS   (MOSELEY) 


H  1 

He    2 

Ne  10 

Li  3 

Na  11 

Gl  4 
Mg  12 

B  5 
Af  13 

C  6 
Si  14 

N  7 
P  15 

O  8 
S  16 

F  9 
Cl  17 

A     18 

K  19 
Cu  29 

Ca  20 
Zn  30 

Sc  21 
Ga  31 

Ti  22 
Ge  32 

V  23 

As  33 

Cr  24 

Se  34 

Mn25 
Br  35 

Fe  26 

Co  27 

Ni  28 

Kr  36 

Rb37 

Ag  47 

Sr  38 
Cd  48 

Y  39 

In  49 

Zr  40 

Sn  50 

Cb  41 
Sb  51 

Mo  42 
Te  52 

—  43 

I  53 

Ru44 

Rh45 

Pd  46 



Xe  54 

Cs  55 
Au  79 

Ba  56 

La  57 

Ce  58 

Ta  73* 

W  74 

—  75 

Os  76 

Ir    77 

Pt   78 

*  The  atomic  numbers  59-72  are  those  of  the  metals  of  the  rare  earths:  Pr  59,  Nd  60, 
-61,  Sa  62,  Eu  63,  Gd  64,  Tb  65,  Dy  66,  Ho  67,  Er  68,  Tm  69,  Yb  70,  Lu  71,  -72. 

It  will  be  seen  that  there  is  a  whole  number  available  for  every 
known  element,  up  to  and  including  gold,  and  not  omitting  the 
rare  elements  which  have  no  satisfactory  place  in  the  periodic 
system.  There  are  two  blank  numbers  in  the  table,  which  corre- 
spond to  two  spaces  below  Mn  in  the  periodic  system,  and  two 
more  amongst  the  rare  elements,  indicating  only  four  elements 
with  atomic  weights  less  than  that  of  gold  yet  to  be  discovered. 
The  atomic  numbers  of  argon  and  potassium  place  them  in  the 
chemically  correct  order,  while  the  atomic  weights  do  not.  The 
same  is  true  of  cobalt  and  nickel  and  of  tellurium  and  iodine. 
Finally,  it  is  evident  that  the  atomic  weight  of  each  element  is, 
roughly,  double  its  atomic  number. 

The  atomic  numbers  represent  the  number  of  unit  positive 
charges  of  electricity  in  the  nucleus  of  the  atom  of  each  element 
(p.  235).  Rutherford  has  shown  that  the  nucleus  contains  almost 
the  whole  mass  of  the  atom,  although  one  or  more  electrons 
(negative)  are  present  also.  Thus,  the  positive  nucleus  of  the 
hydrogen  atom  is  1800  times  heavier  than  one  electron.  The 
nucleus,  however,  is  very  minute,  having  a  diameter  only  about 
one-eighteen  hundredth  of  that  of  an  electron. 

The  atomic  numbers  apparently  determine  all  the  properties 
of  each  element,  and  are  more  fundamental  than  the  atomic 
weights.  The  latter  are  secondary  properties,  in  most  cases 
modified  by  other  factors,  and  in  a  few  cases  actually  thrown  out 
of  order  by  such  factors. 


Crystal  Structure.  —  In  this  connection  it  may  be  mentioned 
that  by  using  crystals  of  different  substances  as  X-ray  gratings, 


THE    PERIODIC    SYSTEM  305 

W.  L.  Bragg  (1914)  has  been  able  to  measure  the  distances  be- 
tween the  rows  of  particles  in  crystals.  He  also  finds  that  the 
particles,  the  regular  arrangement  of  which  gives  the  structure 
(p.  82)  of  the  crystal  (e.g.,  a  cube  of  common  salt),  are  not  the 
molecules  of  the  compound,  much  less  aggregates  of  such  mole- 
cules, but  the  atoms  of  the  constituent  elements.  It  would  thus 
appear  that  the  physical  forces  (if  we  may  call  them  physical) 
which  hold  the  crystalline  solid  together  have  completely  crushed 
the  chemical,  molecular  structure  out  of  existence,  and  have  ar- 
ranged the  constituent  atoms,  as  the  units  of  the  structure,  in  a 
crystallographic  pattern.  Of  course,  when  the  crystal-form  is  de- 
stroyed, by  melting,  solution,  or  vaporization,  the  neighboring 
atoms  remain  united  in  groups,  constituting  the  chemical  mole- 
cules of  the  substance. 

Exercises.  —  1.   Can  you  explain  the  presence  of  free  selenium 
in  the  flues  of  pyrite  burners  (p.  294)? 

2.  How  should  you  attempt  to  obtain  H2Te,  and  what  physical 
and  chemical  properties  should  you  expect  it  to  possess? 

3.  Make  a  list  of  bivalent  elements  and  criticize  this  method  of 
grouping  as  a  means  of  chemical  classification. 

4.  Write  down  the  symbols  of  the  elements  in  the  fourth  series 
(that  beginning  with  rubidium,  and  ending  with  iodine)  on  p.  299. 
Record  the  valence  of  each  element  toward  oxygen,  using  for  refer- 
ence the  chapters  in  which  the  oxygen  compounds  are  described. 


CHAPTER  XXIII 

OXIDES  AND  OXYGEN  ACIDS  OF  THE  HALOGENS 
OXIDATION  AND  REDUCTION 

THE  chief  subjects  of  practical  importance  touched  upon  in  the 
first  part  of  this  chapter  are  connected  with  bleaching  powder 
CaCl(OCl),  and  potassium  chlorate  KC103  and  perchlorate  KC104. 
Hence  our  attention  will  be  largely  directed  to  the  modes  of  making 
these  substances  and  to  their  relations  to  one  another.  Inciden- 
tally, we  shall  encounter  many  actions  of  a  complex  and,  to  us, 
more  or  less  novel  kind. 

Compounds  of  Chlorine  Containing  Oxygen.  —  The  fol- 
lowing are  the  names  and  formulae  of  the  substances: 

HC10  Hypochlorous  acid,  C^O  Hypochlorous  anhydride, 
[HC102]  Chlorous  acid,  ...    .......... 

............  C102  Chlorine  dioxide, 

HC103  Chloric  acid,  .............. 

HC104  Perchloric  acid,  C^Oy  Perchloric  anhydride. 


There  are  also  salts  of  these  acids,  like  the  three  substances 
mentioned  in  the  first  paragraph.  Chlorous  acid  is  itself  unknown, 
but  potassium  chlorite  KC102  and  some  other  derivatives  have 
been  made. 

The  two  anhydrides  (p.  94),  when  brought  into  contact  with 
water,  combine  with  it  to  form  the  acids  opposite  which  they  stand 
in  the  table.  Chlorine  dioxide  (q.v.),  however,  is  not  related  to  any 
one  acid  in  this  way. 

All  these  compounds  differ  from  most  that  we  have  hitherto  dis- 
cussed, inasmuch  as  not  one  of  them  can  be  made  by  direct  union 
of  the  simple  substances. 

Nomenclature  of  the  Acids  and  their  Salts.  —  The  acids 
and  salts  are  named  on  a  plan  similar  to  that  used  in  the  case  of 
the  sulphur  acids: 

306 


OXIDES   AND    OXYGEN   ACIDS   OF   CHLORINE  307 

KC10  Potassium  hypochlorite,  HC10  Hypocldorous  acid, 

KC102  Potassium  chlorite,  HC102  Chlorous  acid, 

KC103  Potassium  chlorate,  HC103  Chloric  acid, 

KC104  Potassium  perchlorate.  HC1O4  Perchloric  acid. 

It  should  be  noted,  however,  that  the  use  of  ic  and  ous  for  more  and 
less  oxygen,  respectively,  and  of  hypo  for  still  less  and  of  per  for 
still  more  oxygen  are  simply  relative  terms  within  a  single  group. 
Thus,  sulphuric  acid  H2S04  has  a  composition  entirely  different 
from  chloric  acid,  and  both  of  these  differ  in  composition  from 
phosphoric  acid  H3PO4.  The  names  and  formulae  of  each  group 
must  be  learned,  separately. 

Chlorine  Monoxide  or  Hypochlorous  Anhydride  ChO.  — 

A  solution  of  pure  hypochlorous  acid  is  most  easily  prepared  by 
dissolving  the  anhydride  in  water.  This  oxide  is  obtained  by 
passing  chlorine  gas  over  warmed  mercuric  oxide  *  HgO  (Fig.  66, 
p.  156).  Each  of  the  constituents  of  the  oxide  combines  with 
chlorine : 

HgO  +  2C12  ->  HgCl2  +  C120. 

The  mercuric  chloride  then  unites  with  another  formula-weight  of 
the  mercuric  oxide  to  form  a  solid  basic  mercuric  chloride  HgO, 
HgCl2,  which  remains  in  the  tube.  The  chlorine  monoxide  is  a 
brownish-yellow,  heavy,  easily  liquefied  gas  (b.-p.  5°).  When 
slightly  warmed  it  decomposes  into  its  constituents  with  explosion. 
The  gas  dissolves  in  water  very  easily  (200  :  1,  by  vol.).  The  yel- 
low solution  of  hypochlorous  acid  which  results: 

C120  +  H20  t=>  2HOC1, 

has  a  strong  odor  of  chlorine  monoxide,  because  the  combination 
is  reversible.  There  are  other  ways  of  preparing  a  dilute  solution 
of  the  acid  (see  below). 

Properties  of  Hypochlorous  Acid.  —  Hypochlorous  acid  is 
unstable,  and  cannot  be  made,  excepting  in  solution,  or  kept,  ex- 

*  The  crystalline,  red  oxide  is  not  sufficiently  active.  The  oxide  must  be 
precipitated  from  sodium  hydroxide  and  mercuric  nitrate  solutions,  it  must 
be  washed  thoroughly  on  a  filter,  and  be  dried  at  300-400°  before  use. 


308  COLLEGE   CHEMISTRY 

cepting  in  dilute  solution.  This  is  in  consequence  of  its  tendency 
to  decompose  in  three  different  ways,  one  of  which,  the  liberation 
of  the  anhydride,  has  just  been  mentioned. 

1.  Hypochlorous  acid  is  a  little-ionized,  weak  add. 


It  neutralizes  active  bases,  its  ionization  equilibrium  being  dis- 
placed forwards  as  the  hydrogen-ion  H+  is  removed  to  form  water: 

NaOH  +  HOC1  ^±  NaOCl  +  H20. 

2.  The  solution,  if  strong,  gives  off  chlorine  monoxide  C120,  the 
union  with  water  being  reversible. 

3.  If  the  solution  is  concentrated,  much  of  the  hypochlorous 
acid  changes  gradually  into  chloric  acid  and  hydrogen  chloride. 
This  is  a  self  -oxidation.    It  occurs  even  in  the  dark: 

3HOC1  -»  HClOa  +  2HC1. 

4.  When  the  solution  is  exposed  to  sunlight,  oxygen  is  evolved. 
rapidly. 

2HOC1  -»  2HC1  +  02. 

This  decomposition  always  takes  place  in  sunlight,  whether  the 
acid  is  present  alone  in  the  water,  or  along  with  other  substances. 
We  have  already  noted  this  fact  in  discussing  chlorine-water  (p. 
162),  which  contains  this  acid. 

5.  In  consequence  of  the  ease  with  which  it  gives  up  oxygen, 
hypochlorous  acid  is  a  strong  oxidizing  agent.     In  this  direction 
it  has  several  important  commercial  applications  (see  below). 

Commercial  Preparation  of  Hypochlorites.  —  For  com- 
mercial purposes,  pure  hypochlorites  are  not,  as  a  rule,  required. 
Hence,  sodium  or  potassium  hypochlorite  is  prepared  by  the  action 
of  sodium  or  potassium  hydroxide  on  chlorine-water.  The  latter 
contains  both  hydrochloric  and  hypochlorous  acids,  and  so  a 
solution  containing  a  mixture  of  sodium  or  potassium  chloride 
and  hypochlorite  is  obtained: 

C12  +  H20  <=±  HC1  +  HOC1.  (1) 

HC1  +  KOH  <=>  KC1  +  H2O.  (2) 

HOC1  +  KOH  <±  KOC1  +  H20.  (3) 


OXIDES   AND   OXYGEN   ACIDS   O^   CHLORINE  309 

Although  action  (1)  is  only  partial,  being  strongly  reversible,  the 
neutralization  of  the  two  acids  in  actions  (2)  and  (3)  displaces  the 
first  equilibrium,  and  all  three  actions  proceed  to  completion. 
Action  (1),  followed  by  (2)  or  (3),  is  a  pair  of  consecutive  actions 
(p.  289),  of  which  the  second  (the  neutralization)  is  the  speedier 
of  the  two.  Both  pairs  of  consecutive  actions  (1)  +  (2)  and  (1) 
+  (3),  can  be  combined  in  one  equation.  Thus,  omitting  the 
water,  which  appears  both  among  products  and  initial  substances 
and  in  any  case  is  present  in  large  excess  as  a  solvent,  and  omitting 
also  the  two  acids,  which  are  used  up  as  quickly  as  they  are  pro- 
.duced  by  equation  (1)  and  are  not  amongst  the  actual  products, 
we  get,  by  addition  of  the  three  equations  (cf.  p.  195),  the  final 
equation: 

Cla  +  2KOH  -» KC1  +  KOC1  +  H20. 

As  lime  is  a  less  expensive  alkali  than  is  potassium  or  sodium 
hydroxide,  it  is  largely  used.  The  chlorine  is  led  into  chambers 
containing  quicklime  CaO  spread  on  trays: 


NOC1 

The  product  is  not  a  mixture,  but  a  mixed  salt  (p.  245),  known 
as  bleaching  powder  or  "  chloride  of  lime."  The  fact  that  this 
is  a  mixed  salt  does  not  interfere  with  its  use  as  a  commercial 
source  of  hypochlorous  acid.  It  is  only  moderately  soluble  in 
water. 

Hypochlorous  Acid  from  Bleaching  Powder.  —  1.  When 
bleaching  powder  is  dissolved  in  water,  being  a  salt,  it  is  very  ex- 
tensively ionized  (see  formulation).  If  now  an  active  acid,  that  is, 
one  giving  a  large  concentration  of  hydrogen-ion,  is  added,  the 
values  of  the  products  of  the  concentrations  [H+]  X  [Cl~]  and 
[H+]  X  [OC1~],  on  which  depend  the  extent  to  which  molecules  of 
HC1  and  HOC1  will  be  formed  (p.  238),  are  large.  HC10,  being 
little  ionized,  is  formed  extensively:  HC1,  being  highly  ionized  is 
formed  in  much  smaller  amount.  Both,  however,  interact  to  pro- 
duce chlorine  and  water,  and  this  displaces  the  other  equilibria. 
Hence  an  active  acid  decomposes  the  salt  almost  completely.  An 


310  COLLEGE    CHEMISTRY 

active  acid  gives,  therefore,  chlorine-water,  and  not  pure  hypo- 
chlorous  acid. 


CaCl(OCl)^Ca++  +  Cr+OCr  .  2«  A  weak   acid> 

H2S04^S04=  +  H+  +  H+  llke    bonc    acid    or    carbonic 

ij         I*  acid,  gives  so  low  a  concen- 

HC1    HOC1  ^rati°n  of  H+  that  union  of 

•  -  T:  -  this  ion  with  OC1~  occurs  to 

*T  form  the  little  ionized  HOC1 

^*2  only,  and  practically  no  com- 

bination of  H+  with  Cl~  takes  place  (see  bleaching). 


CaCl(OCl)  ^  Ca-H-  +  CP  +  OC 
H2C03  <=*  C03=  +  H+  +  H+ 


} 
[ 


When  the  dilute  mixture  is  distilled,  chlorine  monoxide  (2HOC1  <=* 
H20  +  C12O)  passes  over  with  the  steam,  and  so  a  dilute  hypo- 
chlorous  acid  can  be  obtained. 

Hypochlorous  Acid  from  Chlorine-  Water.  —  An  interesting 
way  of  obtaining  dilute  hypochlorous  acid  is  to  add  chalk  CaCO3 
to  chlorine-water  and  distil.  Here,  the  chalk  is  insoluble,  and  so 
gives  a  very  low  concentration  of  Ca++  +  C03=.  The  HC1  in  the 
chlorine-water  gives,  however,  a  sufficiently  large  concentration 
of  H+  to  combine  with  the  CO3=  to  form  H2C03,  which  is  hardly 
ionized  at  all.  This  carbonic  acid  H2CO3  then  decomposes  and 
carbon  dioxide  is  liberated: 


2HC1 

The  hypochlorous  acid,  however,  remains  molecular  HOC1,  gives 
almost  no  H+,  and  so  for  the  most  part  remains  unaffected.  It 
can  afterwards  be  distilled  off  with  the  water. 

Hypochlorous  Acid  as  an  Oxidizing  Agent.  —  Hypochlorous 
acid,  in  decomposing  into  oxygen  and  hydrochloric  acid,  gives 
off  heat.  HOCl,Aq  —  »  HCl,Aq  +  0  +  9300  cal  Hence  more 
energy  is  liberated  in  oxidation  by  the  acid  than  in  oxidation  by 
free  oxygen,  and  the  former  is  therefore  more  active  as  an  oxidizing 
agent  (p.  224).  Thus,  hypochlorous  acid,  either  in  pure  solution 
or  in  the  form  of  chlorine-water,  oxidizes  sulphurous  acid  instantly  : 
H2S03  +  HOC1  -»  H2S04  +  HC1. 


OXIDES   AND    OXYGEN   ACIDS   OF   CHLORINE  311 

It  also  oxidizes  bromine  and  iodine,  in  water,  although  these  ele- 
ments are  not  affected  by  free  oxygen,  giving  bromic  and  iodic 
acids,  respectively: 

5HC10  +  I2  +  H20  -*  5HC1  +  2HIO3. 

The  solution  also  oxidizes  organic  colored  substances  (p.  221), 
producing  colorless,  or  less  strongly  colored  ones.  Thus,  it 
oxidizes  indigo  (deep  blue)  quickly  to  isatin,  a  yellow  substance 
relatively  pale  in  color: 

Ci6H10N202  +  2HOC1  ->  2C8H5N02  +  2HC1. 

In  ways  just  as  definite  as  this,  hypochlorous  acid  will  change  the 
composition  of  other  colored  substances,  although,  since  we  do  not 
know  the  formulae  of  all  these  substances,  we  cannot  always  write 
equations  for  the  actions. 

Hypochlorous  Acid  as  a  Bleaching  Agent.  —  It  is  on  account 
of  its  oxidizing  power  that  hypochlorous  acid  is  used  commercially 
in  bleaching.  It  is  not  applied  to  paints,  which  are  chiefly  mineral 
substances,  but  to  complex  compounds  of  carbon,  such  as  consti- 
tute the  coloring  matters  of  plants  and  of  those  artificial  dyes 
which  are  now  manufactured  in  great  variety. 

Cotton  and  linen,  in  their  original  states,  are  not  pure  white. 
Bleaching  is,  therefore,  an  extensive  and  most  important  industry. 
The  yarn  or  cloth  must  first  be  freed  from  cotton-wax  and  tannin, 
since  the  former  would  protect  it  from  the  action  of  the  bleaching 
agent,  and  both  would  make  the  subsequent  dyeing  uneven.  The 
material  is,  therefore,  first  boiled  with  very  dilute  sodium  hydroxide 
solution,  and  washed  with  water.  The  goods  are  then  saturated 
with  bleaching  powder  solution,  and  piled  loosely  until  the  coloring 
matter  has  been  oxidized.  They  are  finally  washed  with  extreme 
thoroughness. 

As  a  rule,  an  active  acid  is  not  added.  The  bleaching  is  pro- 
duced by  the  hypochlorous  acid  liberated  by  the  action  of  the  car- 
bon dioxide  from  the  air.  The  carbon  dioxide  dissolves  in  the 
water  of  the  solution  on  the  goods,  and  forms  carbonic  acid: 
C02  +  H20  <=»  H2C03  (see  p.  310,  par.  2).  The  subsequent  wash- 
ing removes  all  traces  of  the  bleaching  powder,  of  the  lime  which 
the  powder  often  contains,  and  of  the  hypochlorous  acid,  which 


312  COLLEGE    CHEMISTRY 

otherwise  would  act  gradually  upon  the  cotton  or  linen  and 
"rot"  it.  Bleaching  agents,  when  used  in  the  household  with- 
out sufficiently  careful  washing,  are  liable  to  cause  serious  damage 
from  this  cause. 

Cotton  and  linen  are  composed  of  cellulose  (CeHuA)*,  a  rather 
inert  substance,  and  one  which  is  very  slowly  acted  upon  by  dilute 
hypochlorous  acid.  Hence,  with  brief  contact  and  proper  han- 
dling, no  damage  is  done.  Wool,  silk,  and  feathers,  however,  are 
composed  largely  of  compounds  (proteins)  containing  nitrogen 
(up  to  15  per  cent)  in  addition  to  the  above  three  elements.  Their 
constituent  material  interacts  as  easily  with  hypochlorous  acid  as 
do  the  traces  of  coloring  substances.  Hence,  since  the  fabric  itself 
would  be  attacked  by  this  agent,  sulphur  dioxide  or  sulphurous 
acid  (p.  289)  is  used  for  bleaching  these  materials. 

Bleaching  Powder  in  Sanitation.  —  A  disinfectant  is  a  sub- 
stance which  destroys  bacteria  and  other  minute  organisms. 
Bleaching  powder  has  a  distinct  odor  of  chlorine  monoxide  (not 
chlorine).  This  is  due  to  the  action  of  atmospheric  carbon  di- 
oxide liberating  hypochlorous  acid  (p.  310).  The  dry  powder 
therefore  will  disinfect  the  air  and  surrounding  objects.  It  must 
be  used  with  discretion,  however,  as  the  gas  is  very  corrosive. 

As  already  mentioned  (p.  91),  in  the  purification  of  city  waters 
the  organisms  which  give  rise  to  typhoid  fever  are  destroyed  by 
adding  a  small  proportion  of  bleaching  powder  solution  (about  20 
Ibs.  per  million  gallons  of  water).  The  salt  is  hydrolyzed  (p.  197), 
giving  a  basic  calcium  chloride  and  free  hypochlorous  acid.  The 
latter  kills  the  organisms,  and  is  itself  decomposed  in  the  process, 
so  that  nothing  offensive  remains  in  the  water.  There  is  only  a 
minute  increase  in  the  proportion  of  salts  of  calcium  (hardness). 

Recently,  chlorine-water,  made  by  use  of  cylinders  of  liquid 
chlorine  (p.  160),  has  in  many  cases  taken  the  place  of  bleaching 
powder  solution  for  this  purpose. 

Chlorine  not  a  Bleaching  Agent.  —  Chlorine  itself  is  often, 
erroneously,  described  as  a  bleaching  agent.  If  a  dry,  colored 
cloth  be  hung  for  a  week  in  chlorine  gas,  dried  by  a  little  sulphuric 
acid  in  the  bottom  of  the  bottle  (Fig.  89),  little  or  no  change  in  the 
color  will  occur.  But  a  wet  rag  is  bleached  as  soon  as  the  chlorine 


OXIDES   AND    OXYGEN   ACIDS   OF   CHLORINE 


313 


has  time  to  dissolve  in  the  water  and  give  the  necessary  hypochlo- 
rous  acid.  Flowers  are  bleached  by  dry  chlorine  gas,  because  by 
their  nature  they  contain  the  indispensable  water. 


Chemical    Properties    of    Hypochlorites.  - 

When  hypochlorites  are  heated  they  change  into 
chlorates  (see  below).  They  may  also  give  off 
oxygen,  2CaCl(OCl)  ->  2CaCl2  +  O2.  Although 
this  decomposition  is  slow  in  cold  solutions  of 
hypochlorites,  or  when  they  are  preserved  in  the 
dry  form,  it  may  be  hastened  by  means  of  cata- 
lytic agents.  The  addition  of  a  little  cobalt  hy- 
droxide (q.v.)  to  a  paste  of  bleaching  powder  and 
water  causes  rapid  evolution  of  oxygen. 


TT 


FIG.  89. 


Chlorates.  —  Like  hypochlorous  acid  itself,  the  hypochlorites 
turn  into  chlorates.  Thus,  when  chlorine  is  passed  into  a  warm, 
concentrated  solution  of  potassium  hydroxide,  and  particularly 
when  an  excess  of  chlorine  is  used,  the  potassium  hypochlorite 
changes  into  potassium  chlorate  KC1O3  as  fast  as  it  is  formed. 
Since  this  action  (equation  2)  requires  3KC10,  the  equation 
formerly  given  (p.  309)  must  be  tripled: 

3C12  +  6KOH  ->  3KC1    +  3KC1O  +  3H2O.  (1) 

3KC1O^2KC1   +  KC1Q3.  (2) 

Adding:      3C12  +  6KOH  ->  KC1O3  +  5KC1  +  3H2O. 

When  the  solution  is  cooled,  the  less  soluble  chlorate  crystallizes. 

This  action  involves  converting  five-sixths  of  the  valuable  potas- 
sium hydroxide  into  the  relatively  less  valuable  potassium  chloride. 
Hence,  in  practice,  the  makers  carry  out  the  corresponding  action 
with  calcium  hydroxide.  They  then  add  potassium  chloride  to  the 
resulting  solution,  containing  calcium  chloride  (very  soluble)  and 
calcium  chlorate  Ca(C103)2.  The  potassium  chlorate,  formed 
by  double  decomposition,  crystallizes  when  the  solution  is 
cooled. 

All  chlorates  are  at  least  moderately  soluble  in  water  (see  Table 
inside  of  front  cover) .  Potassium  chlorate  is  used  in  making  fire- 
works, explosives,  and  matches.  An  intimate  mixture  with  sugar 
Ci2H22On  burns  with  semi-explosive  violence,  the  oxygen  of  the 


314  COLLEGE    CHEMISTRY 

salt  combining  with  the  carbon  and  hydrogen  of  the  sugar  to  form, 
carbon  dioxide  and  water. 

Chloric  Acid  HCIO3.  —  Since  none  of  the  acids  of  this  series 
can  be  obtained  by  direct  union  of  their  elements  (p.  306),  it  is 
usual  first  to  prepare  the  salts,  and  to  make  the  acids  from  the 
salts  by  double  decomposition.  This  acid  may  be  obtained,  in 
solution  in  water,  by  adding  the  calculated  amount  of  diluted 
sulphuric  acid  to  a  solution  of  barium  chlorate: 

Ba(C103)2  +  H2S04  fc?  BaS04|  +  2HC103. 

The  barium  sulphate,  being  insoluble,  is  removed  by  filtration. 
It  will  be  noted  that  double  decomposition  involving  precipitation 
may  thus  be  used  for  obtaining  a  soluble  product,  as  well  as  an  in- 
soluble one  (cf.  selenic  acid,  p.  294). 

The  solution  may  be  concentrated  (to  about  40  per  cent)  by 
evaporation,  but  must  not  be  heated  above  40°,  as  the  acid  decom- 
poses near  this  temperature.  The  resulting  thick,  colorless  liquid 
has  powerful  oxidizing  qualities,  setting  fire  to  paper  (made  of 
cellulose  (C6Hi006)i)  which  has  been  dipped  into  it.  It  converts 
iodine  into  iodic  acid,  2HC103  +  I2  -*  2HIO3  -f  C12.  When 
not  in  solution,  or  when  warmed  in  solution  beyond  40°,  the  acid 
decomposes,  giving  chlorine  dioxide  and  perchloric  acid: 

3HC1O3  -»  H20  +.2C1Q2  +  HC104. 

Chlorine  Dioxide:  Chlorous  Acid.  —  Chlorine  dioxide  C102 
(see  above)  is  a  yellow  gas  which  may  be  liquefied,  and  boils  at 
+  10°.  The  gas  and  liquid  are  violently  explosive,  the  substance 
being  resolved  into  its  elements  with  liberation  of  much  heat.  It 
is  formed  whenever  chloric  acid  is  set  free,  and  hence  it  is  seen 
when  a  little  powdered  potassium  chlorate  is  touched  with  a  drop 
of  concentrated  sulphuric  acid  (end  of  last  section)*.  Concen- 
trated hydrochloric  acid  turns  yellow  from  the  same  cause  when 
any  chlorate  is  added  to  it.  These  actions  are  used  as  tests  for 
chlorates,  and  distinguish  them  from  perchlorates  (q.v.).  With 

*  The  mixture  of  sugar  and  potassium  chlorate  (p.  313)  can  be  set  on  fire 
by  a  drop  of  sulphuric  acid.  The  latter  liberates  chloric  acid,  which  in  turn 
gives  C1O2,  and  the  latter,  being  a  violent  oxidizing  agent,  starts  the  combus- 
tion of  the  sugar. 


OXIDES  AND    OXYGEN   ACIDS   OF   CHLORINE  315 

water,  chlorine  dioxide  gives  a  mixture  of  chlorous  acid  HC102 
and  chloric  acid,  and  with  bases  a  mixture  of  the  chlorite  and 
chlorate. 

Perchlorates.  —  When    heated,    chlorates    give    perchlorates. 

Chlorates  also  give  oxygen  at  the  same  time  (p.  27) : 

(2KC1O3-»2KC1  +  3O2, 
(4KC103  ->  3KC104  +  KCL 

These  actions,  like  the  three  decompositions  of  hypochlorous 
acid  (p.  308),  are  independent,  and  proceed  simultaneously. 
They  are  concurrent  reactions  (see  below).  Their  relative  speed, 
however,  varies  with  the  temperature,  and  the  decomposition  into 
chloride  and  oxygen  may  completely  outrun  the  other  when  a 
catalytic  agent  like  manganese  dioxide  is  added  (p.  29).  When 
pure  potassium  chlorate  is  heated  cautiously,  about  one-fifth  of 
it  has  lost  all  its  oxygen  by  the  time  the  rest  has  turned  into  per- 
chlorate.  The  mixture  may  be  separated  by  grinding  with  the 
minimum  quantity  of  water  which  will  dissolve  the  chloride  it  con- 
tains. The  perchlorate,  having  at  15°  less  than  one-twentieth  of 
the  solubility  of  the  chloride,  will  remain,  for  the  most  part,  un- 
dissolved.  The  perchlorates  are  much  more  stable  (p.  93)  than 
the  chlorates,  or  hypochlorites :  they  are  all  soluble  in  water,  and 
they  are  used  in  making  matches  and  fireworks. 

Perchloric  Acid  HCIO^  and  Perchloric  Anhydride  C12O7.  — 

Pure  perchloric  acid  explodes  when  heated  above  92°.  But,  like 
other  liquids,  its  boiling-point  is  lower  when  its  vapor  is  under 
reduced  pressure  (cf.  p.  87).  At  56  mm.  pressure  it  boils  at  39°,  a 
temperature  at  which  hardly  any  decomposition  is  noticeable. 
Hence  the  acid  may  be  made  by  mixing  potassium  perchlorate 
and  concentrated  sulphuric  acid  and  distilling  the  mixture  cau- 
tiously in  a  vacuum  (p.  222) : 

KC104  +  H2SO4  *=?  KHS04  +  HC104T. 

Perchloric  acid  is  a  colorless  liquid,  which  decomposes,  and  often 
explodes  spontaneously,  when  kept.  A  70  per  cent  solution  in 
water  is  perfectly  stable,  however.  Although  it  is  an  active  oxidiz- 
ing agent,  it  is  not  so  active  as  chloric  acid,  and  does  not  oxidize 


316  COLLEGE    CHEMISTRY 

hydrogen  chloride  in  cold  aqueous  solution.  Hence  a  drop  of 
hydrochloric  acid  placed  on  a  crystal  of  a  perchlorate  gives  no 
yellow  color.  When  the  acid  is  liberated  by  concentrated  sulphuric 
acid,  it  does  not  at  once  give  the  yellow  chlorine  dioxide  (p.  314). 
Perchloric  anhydride  C12O7  may  be  prepared  by  adding  phosphoric 
anhydride  to  perchloric  acid  in  a  vessel  immersed  in  a  freezing  mix- 
ture, P206  +  2HC104  — >  2HP03  +  C1207.  Phosphoric  anhydride 
is  often  used  in  this  way  for  removing  the  elements  of  water  from 
compounds.  It  combines  with  the  water  to  form  metaphosphoric 
acid  HP03.  By  gently  warming  the  mixture,  the  perchloric 
anhydride  can  be  distilled  off.  It  is  a  colorless  liquid  boiling  at 
82°  (760  mm.)  and  exploding  when  struck  or  too  strongly  heated. 

Relation  of  Anhydride  and  Acid  or  Salt.  —  The  derivation 
of  the  formula  of  the  anhydride  from  that  of  the  acid  or  salt 
should  receive  special  attention.  '  In  the  mind  of  the  chemist,  the 
one  always  instantly  suggests  the  other,  so  often  does  he  think 
of  them  as  potentially  the  same  substance.  The  beginner,  how- 
ever, finds  this  habit  hard  to  acquire,  and  indeed  is  more  likely  to 
blunder,  in  trying  to  divide  the  formula  of  an  acid  into  the  formula? 
of  water  and  the  anhydride,  than  in  any  other  calculation  he  makes. 

The  rule  is :  If  the  formula  of  the  acid  shows  an  even  number 
of  hydrogen  atoms  (H2S04  or  EUSiC^),  subtract  all  the  elements 
of  water  (H20  or  2H2O),  and  the  balance  is  the  anhydride  (SO3 
or  Si02).  The  divided  formulae  are  H20,SO3  or  2H20,Si02.  If 
there  is  an  odd  number  of  hydrogen  atoms  (HC1O4  or  H3P04) 
double  the  formula  (H2C1208  or  H6P208),  and  subtract  all  the  ele- 
ments of  water  as  before  (C12O7  or  P2O5).  Then  check  the  result, 
by  adding  the  water  again,  and  dividing  by  two,  correcting  the 
blunder  if  one  has  been  made. 

If  the  substance  is  a  salt  (CuSO4  or  KC104),  subtract  the  oxide 
of  the  metal  (CuO  or  K20),  taking  care  to  assign  to  the  metal  the 
same  valence  in  the  oxide  as  it  shows  in  the  salt  (S03  or  C1207). 

There  are  several  uses  for  this  art  of  ascertaining  the  anhydride 
corresponding  to  a  given  salt  or  acid.  One  is  in  the  making  of 
equations  (see  p.  325).  Another  is  in  finding  the  valence  of 
the  non-metal.  Thus,  in  KC104  the  anhydride  is  C1207,  and  the 
valence  of  the  chlorine  is  seven.  In  H3PO4  the  anhydride  is  P2O5 
and  the  phosphorus  quinquivalent.  In  HP03  (metaphosphoric 


OXIDES   AND    OXYGEN   ACIDS    OF   CHLORINE  317 

acid),  the  anhydride  is  again  P2O5,  and  the  phosphorus  is  therefore 
in  the  same  state  of  oxidation  —  both  are  phosphoric  acids. 

Simultaneous,  Independent  Chemical  Changes  in  the 
Same  Substances.  —  When  two  or  more  reactions  go  on  simul- 
taneously in  the  same  materials,  the  actions  may  be  consecutive 
(p.  289)  or  they  may  be  parallel.  In  the  latter  case  they  are  called 
concurrent  reactions.  Thus,  hypochlorous  acid  undergoes  three 
different  changes: 

2HC10  ->  H20  +  C120. 

3HC10  -»  HC103  +  2HC1. 

2HC10  ->  2HC1  +  O2. 

Some  molecules  decompose  into  water  and  chlorine  monoxide  (p. 
308),  while  others  give  chloric  acid  and  hydrogen  chloride,  and  still 
others  hydrogen  chloride  and  oxygen.  Since  the  same  molecule 
cannot  undergo  more  than  one  of  these  different  changes,  it  follows 
that  the  actions  are  independent  of  one  another.  This  is  shown 
by  the  fact  that  in  sunlight  the  third  predominates,  while  in  the 
dark  it  falls  far  behind  the  second.  Since  the  relative  quantities 
of  the  products  vary,  the  several  simultaneous  actions  cannot  be  put 
in  the  same  equation.  The  fundamental  property  of  an  equation  is 
to  show  the  constant  proportions  by  weight  between  every  pair  of 
substances  in  it.  Hence  three  separate  equations  are  required  in 
the  present,  and  in  all  similar  cases  where  all  the  proportions  are 
not  constant.  Thus,  again,  in  the  decomposition  of  potassium 
chlorate  by  heating  (p.  315),  it  would  be  misleading  and  wrong  to 
add  the  two  equations  together  and  write,  for  the  whole  action: 

2KC103  ->  KC1  +  KC104  +  O2. 

This  equation  would  mean  that  the  proportions  amongst  the  prod- 
ucts were  always  KC1  :  KC1O4  :  O2  or  74.6  :  138.6  :  32,  whereas, 
in  fact,  the  proportions  vary  with  the  conditions  —  the  tempera- 
ture used  or  the  presence  of  a  catalyst  which  hastens  one  action 
but  not  the  other. 

Consecutive  reactions  (p.  289),  however,  like  (1)  followed  by  (2) 
on  pp.  308,  313,  may  be  combined  in  one  equation,  since  in  them 
all  the  proportions  must  necessarily  be  constant.  These  equations 
are  interlocked,  for  (2)  consumes  what  (1)  produces. 


318  COLLEGE    CHEMISTRY 

Oxygen  Acids  of  Bromine.  —  No  oxides  of  bromine  have  been 
made,  but  the  acids  HBrO  (hypobromous  acid)  and  HBr03  (bromic 
acid)  and  their  salts  are  familiar. 

By  the  action  of  bromine  on  dilute,  cold  potassium  hydroxide 
solution,  potassium  bromide  and  hypobromite  are  formed: 

Br2  +  2KOH  ->  KBr  +  KBrO  +  H2O. 

When  the  solution  is  heated,  the  hypobromite  turns  into  potassium 
bromate  and  bromide.  The  actions  are  exact  parallels  of  the  cor- 
responding ones  for  chlorine  (pp.  309,  313). 

Aqueous  bromic  acid  HBrO3  may  be  made  in  the  same  way  as 
chloric  acid  (p.  314),  or  by  the  action  of  chlorine-  water  on  bromine: 

5HC10  +  H20  +  Br2  ->  2HBr03  +  5HC1. 

The  solution  is  colorless  and  has  powerful  oxidizing  properties. 
Thus,  it  converts  iodine  into  iodic  acid  :  2HBr03  +  12  —  >  2HIO3  +  Br2. 
It  appears,  therefore,  that  iodine  has  more  affinity  for  oxygen  than 
has  bromine. 

The  Oxide  and  Oxygen  Acids  of  Iodine.  —  The  following  are 
the  familiar  acids  and  their  corresponding  salts: 

HI03  Iodic  acid,  KI03         Potassium  iodate, 

[HIC>4  Periodic  acid],  NaIC>4       Sodium  periodate, 

H5IOe  Periodic  acid,  NasHslOe  Disodium  periodate. 


There  is  one  oxide,  iodic  anhydride  I205. 

Sodium  Iodate  NaIO3  is  found  in  Chile  saltpeter.  It  may 
be  made,  in  much  the  same  fashion  as  are  the  chlorates  and 
bromates  (pp.  313,  318),  by  adding  powdered  iodine  to  a  hot  solu- 
tion of  potassium  or  sodium  hydroxide.  It  is  disodium  periodate 
Na^HsIOe,  however,  which,  being  least  soluble,  crystallizes  out. 

Iodic  Acid  HI03  is  formed  by  passing  chlorine  through  iodine 
suspended  in  water.  The  action  is  parallel  to  that  of  chlorine  on 
bromine-water: 

5HC10  +  H20  +  I2  -*  2HI03  +  5HC1. 

A  still  better  way  is  to  boil  iodine  with  aqueous  nitric  acid  (q.v.). 
The  latter  gives  up  oxygen  readily,  and  is  here  used  solely  on  this 


OXIDES  AND   OXYGEN  ACIDS,    THE   HALOGENS  319 

account.  Hence,  it  may  be  omitted  from  the  equation,  only  the 
oxygen,  of  which  it  is  the  source,  appearing: 

I2  +  H20  +  50  ->  2HI03. 

In  both  these  actions  the  initial  substances  (including  the  excess  of 
nitric  acid)  and  the  products,  with  the  exception  of  the  iodic  acid 
itself,  are  all  volatile.  When  the  solution  is  concentrated  by  evap- 
oration, therefore,  only  the  iodic  acid  crystallizes.  It  is  a  white 
solid,  perfectly  stable  at  ordinary  temperatures,  and  can  be  kept 
indefinitely.  At  170°  it  begins  to  give  off  water  vapor  (2HIOS  +± 
H2O  +  I205),  leaving  iodic  anhydride.  The  latter  is  a  white 
crystalline  powder  which  may  be  raised  to  300°  before  it,  in  turn, 
breaks  up,  giving  iodine  and  oxygen. 

Chemical  Relations.  —  The  compounds  of  the  halogens  with 
metals  and  with  hydrogen  diminish  in  stability,  with  ascending 
atomic  weight  of  the  halogen,  in  the  order:  F  (19),  Cl  (35.5),  Br  (80), 
I  (127).  Each  halogen  will  displace  those  following  it  from  this 
kind  of  combination.  In  the  case  of  the  oxygen  compounds,  the 
order  of  stability  is  just  the  reverse,  those  of  iodine,  for  example, 
being  the  only  ones  which  are  reasonably  stable. 

Amongst  the  oxygen  acids  of  any  one  halogen,  those  containing 
most  oxygen  are  most  stable.  The  salts  are  in  all  cases  more  stable 
by  far  than  the  corresponding  acids. 

The  halogens  when  combined  with  metals  and  hydrogen  are 
univalent  (HI,  KC1,  etc.).  It  is  clear,  however,  that,  when  united 
with  oxygen,  their  valence  is  higher.  The  maximum  is  shown  in 
perchloric  anhydride  (C12O7),  where  chlorine  appears  to  be  septi- 
valent. 

The  formulae  of  the  acids  might  be  written  so  as  to  retain  the 
univalence : 

H-C1,  H-O-C1,  H-O-O-C1,  H-O-0-0-C1, 
H-O-O-O-O-C1. 

But  compounds  in  which  we  are  compelled  to  believe  that  two  oxy- 
gen units  are  united  are  usually  unstable  (e.g.,  hydrogen  peroxide, 
H  — O  — O  — H),  and  we  should  expect  the  instability  would  be 
greater  with  three  and  with  four  units  of  oxygen  in  combination. 
Here,  however,  the  reverse  state  of  affairs  must  be  taken  account 


320  COLLEGE    CHEMISTRY 

of  in  our  formulae,  for  HC1O4  is  the  most  stable  of  the  chlorine  set. 
This  reasoning,  together  with  the  septivalence  in  C^O?,  leads  us 
to  assume  the  valence  seven  in  perchloric  acid  (see  Periodic  system) 
The  structural  formulae  (cf.  p.  292)  of  some  of  these  substances  are 
therefore  written  as  follows: 

O  O 

II  II 

H-C1,    H-O-C1,    H-O-C1  =  O,    Na-O-I  =  O. 

II  II 

O  O 

OXIDATION  AND  REDUCTION 

Oxidation  by  Oxygen.  —  The  simplest  oxidations  are  the 
cases  where  a  metal  or  non-metal  unites  with  oxygen: 

2Cu  +  02  ->  2CuO,        S  +  02  -»  S02. 

Union  of  a  compound  with  additional  oxygen  is  oxidation  also. 
2S02  +  02  -» 2S03,        3KC10  -»  2KC1  +  KC103. 

The  removal  of  hydrogen  from  hydrogen  chloride  (preparation  of 
chlorine,  p.  156),  is  also  denned  as  oxidation. 

02  +  4HC1  ->  2H20  +  2C12. 
2KMn04  +  16HC1  -» 8H2O  +  2KC1  +  2MnCl2  +  5C12. 

Every  oxidation  is  accompanied  by  reduction  of  the  oxidizing  agent. 
Thus,  in  the  second  last  equation,  the  free  oxygen  is  reduced  to 
water.  Again,  in  the  third  last  equation,  2KC10  is  reduced  to 
2KC1,  while  1KC1O  becomes  KC1O3  by  oxidation. 

In  the  laboratory,  we  frequently  discover  that  an  oxidation  has 
occurred  by  noticing  the  presence  of  a  product  of  reduction. 
Thus,  when  we  heat  carbon  with  sulphuric  acid:  2H2S04  +  C  — > 
C02  +  2H20  +  2S02,  we  do  not  notice  the  product  of  oxidation, 
C02,  because  it  is  odorless  and  colorless,  but  we  perceive  at  once 
the  odor  of  the  sulphur  dioxide,  and  realize  that  the  sulphuric  acid 
must  have  oxidized  some  substance,  or  this  gas  would  not  have 
been  formed  at  the  temperature  employed. 

Note  that  the  removal  of  the  elements  of  water  is  neither  oxida- 
tion nor  reduction,  for  equivalent  amounts  of  both  oxygen  and 
hydrogen  are  removed: 

2HC10  ->  H20  +  ClaO,        H2CO3  -»  H20  +  C02. 


OXIDATION   AND    REDUCTION  321 

In  the  cases  discussed  above,  oxidation  consists  always  in  adding 
oxygen  or  removing  hydrogen. 

Oxidation  by  Other  Negative  Elements.  —  Oxygen  is  only 
one  of  the  class  of  elements  called  non-metallic  or  negative  ele- 
ments, so  we  cannot  logically  restrict  the  term  " oxidation"  to 
actions  involving  oxygen.  Thus,  forming  a  chloride,  or  increasing 
the  proportion  of  chlorine  in  a  compound  is  oxidation: 
Cu  +  C12  -*  CuCl2,  2FeCl2  +  C12  -*  2FeCl3. 

In  every  compound,  one  of  the  elements  is  relatively  positive 
and  the  other  relatively  negative.  Thus,  copper  is  positive  and 
chlorine  negative.  In  carbon  dioxide  C02,  carbon  is  (relatively) 
positive  and  oxygen  negative,  and  in  calcium  carbide,  CaC2;  cal- 
cium is  positive  and  carbon  (relatively)  negative. 

Thus,  oxidation  is  introducing,  or  increasing  the  proportion  of  the 
negative  element,  or  removing,  or  reducing  the  proportion  of  the 
positive  element.  Reduction  is  the  converse. 

Oxidation  and  Valence.  —  Combining  a  metal  with  oxygen 
or  sulphur  raises  the  active  valence  of  the  metal  from  zero  to  some 
finite  value:  2Cu°  +  02°  ->  2CuII011.  Metallic  copper  has  no 
valence  in  use.  In  CuO  or  CuCl2  it  has  gained  the  valence  II. 
The  copper  has  been  oxidized.  Similarly,  changing  FeCl2  into 
FeCl3  increases  the  active  valence  of  the  iron  from  II  to  III  (oxida- 
tion). Conversely,  changing  2HC1  to  C12  decreases  the  active 
valence  of  chlorine  from  I  to  zero  (oxidation).  In  the  same 
equation  (p.  320),  KMn  in  KMnCX  must  have  a  total  valence  of 
VIII,  but  in  the  products  KC1  +  MnCl2  the  total  valence  has 
decreased  to  III  (reduction). 

Again,  in  displacement,  e.g.,  Zn  +  2HC1  — >  ZnCl2  +  H2,  the 
zinc  is  oxidized  because  the  active  valence  goes  from  zero  to  II, 
and  the  hydrogen  is  reduced. 

Hence,  oxidation  consists  in  increasing  the  active  Valence  of  a 
positive  element  or  decreasing  that  of  a  negative  element.  Reduc- 
tion is  the  converse. 

This  way  of  stating  the  rule  makes  it  clear  why  removing  the 
elements  of  water  is  neither  oxidation  nor  reduction.  We  are 
removing  both  a  positive  and  a  negative  element,  and  are  removing 
them  in  equi-valent  amounts,  2H1  +  O11. 


322  COLLEGE    CHEMISTRY 

Oxidation  and  lonization.  —  If,  in  the  last  illustration,  we 
write  the  equation  ionically :  Zn  +  2H+  — >  Zn++  +  H2,  we  dis- 
cover that,  logically,  we  must  consider  the  change  from  metallic 
zinc  to  zinc-ion  to  be  in  itself  oxidation.  This  is  the  case  whether 
the  zinc-ion  later  combines  with  a  negative  ion  to  form  a  molecule 
or  not.  Mere  union  or  disunion  of  ions  is  neither  oxidation  nor 
reduction.  Conversely,  the  discharge  of  the  2H+  giving  H2  is 
reduction. 

Thus,  ionization  of  an  elementary  substance  to  form  a  positive 
ion  is  oxidation,  and  ionization  to  form  a  negative  ion  is  reduction, 
and  conversely. 

Oxidation  and  Electrons.  —  Increasing  the  valence  of  an 
atom  of  a  positive  element  (oxidation)  consists  in  removing  one  or 
more  electrons :  Na°  —  e  =  Na+  (p.  235) .  Increasing  the  valence 
of  an  atom  of  a  negative  element  (reduction)  means  adding  one  or 
more  electrons :  Cl°  +  e  — >  Cl~. 

Hence,  oxidation  is  removing  electrons  and  reduction  is  adding 
electrons. 

Making  Equations  for  Oxidations  and  Reductions.  —  The 

writing  of  equations  for  actions  involving  oxidation  and  reduction, 
where  there  are  more  than  two  substances  on  one  side  of  the  equation, 
is  difficult,  and  a  system  or  plan  is  of  great  value.  The  plan  of 
partial  equations  (p.  194)  is  often  helpful.  There  are  three  other 
systems  which  are  in  use.  (1)  When  the  action  involves  oxygen 
acids  and  their  salts,  the  formulae  can  be  rewritten  so  as  to  show 
the  anhydride  (see  below).  (2)  The  second,  called  the  system  of 
positive  and  negative  valences,  is  more  generally  applicable  (next 
section).  (3)  The  third  describes  oxidation  in  terms  of  ions  and 
positive  electrical  charges  (p.  325). 

Making  Equations:  Using  Positive  and  Negative  Va- 
lences (p.  276).  —  1.  Each  compound  is  composed  of  elements 
which  are,  relatively  to  one  another,  either  positive  or  negative. 
Thus,  in  KMn04,  K  and  Mn  are  positive  and  0  is  negative.  In 
CS2,  C  is  (relatively)  positive  and  S  negative  (see  p.  277).  We 
say,  then,  that  C  has  a  positive  valence  of  four  (+4)  and  S  has  a 
negative  valence  of  two  (  —  2),  just  as  it  has  in  H2S. 


OXIDATION   AND   REDUCTION  323 

2.  In  each  compound,  the  algebraic  sum  of  the  positive  and  negative 
valences  must  be  zero.     Thus,  in  CS2  the  sum  is +4  —  2x2  =  0 
(CttS2=).     This  is  simply  the  rule  of  equi-valence  (p.  62),  with  the 
addition  of  the  idea  of  relative  positiveness  and  negativeness. 

This  enables  us  to  determine  the  valence  of  each  element  in  a 
compound  like  KMnO4.  K+  is  always  univalent  and  positive.  0=, 
in  inorganic  compounds,  is  always  bivalent  and  negative.  The 
valence  of  Mn  has  different  values:  MnnCl2,  Mn2m03,  Mn^Ou, 
Mn2VII07,  etc.  By  the  rule  (sum  of  valences  equals  zero)  we  can 
tell  the  valence  of  Mn  in  this  compound.  The  valence  of  64 
(40=)  is  -8.  That  of  K  is  +1.  That  of  Mn  must  therefore  be 
+7  (KMn+vn04).*  Again,  in  HC103,  the  valence  of  03  is  -6, 
that  of  H  is  +1,  therefore  that  of  Cl  must  be  +5.  Still  again,  in 
K2Cr2O7,  the  valence  of  O7  is  -14,  that  of  K2  is  +2,  that  of  Cr2 
is  therefore  +12,  and  that  of  Cr  necessarily  +6  (K2Cr2+VI07). 

3.  Since  rule  2  applies  to  every  compound  used  or  produced  in 
a  chemical  change,  it  follows  that  when  in  a  reaction  the  valence  of 
an  element  changes  in  value,  that  of  one  or  more  of  the  other  elements 
must  also  change,  so  as  to  maintain  the  equality  of  +  and  —  valences. 
Thus,  if  one  element  loses  in  valence,  to  the  extent  of  +6,  some 
other  element  (or  elements)  must  lose  —6,  or  gain  +6.     The  gain 
(or  loss)  of  one  element  must  cancel  the  gain  (or  loss)  of  some  other 
element. 

4.  The  valence  of  a  free  element,  that  is,  its  active  valence,  is 
zero.     A  free  element  is  also  neutral  —  neither  positive  nor  nega- 
tive —  because  it  is  not  combined  with  any  other  element. 

Illustration  of  rules  4  and  5.  Thus,  in  the  action  for  preparing 
chlorine  with  manganese  dioxide  (p.  158) : 

Mn02  +  4HC1  -» MnCl2  +  2H20  +  C12, 

4H  (4H+)  has  the  valence  +4  on  both  sides.  On  the  left  side, 
4C1  (4C1~)  has  the  valence  —4:  on  the  right,  2C1  has  the  valence 
—2,  and  C12  has  the  valence  0.  So  far  as  chlorine  is  concerned, 
there  is  a  change  from  —4  to  —2,  or  a  difference  of  —2.  Again, 
on  the  right,  Mn  has  the  valence  +2,  while  on  the  left  side 
it  has  the  valence  +4,  a  difference  of  +2.  The  two  differences, 
—  2  and  +2,  cancel  one  another.  Stated  otherwise,  manganese 

*  The  reader  should  write  this,  and  other  formulae  discussed  below,  so  as  to 
show  the  valences  thus:  K+Mnfl^+Or  (cf.  p.  276). 


324  COLLEGE    CHEMISTRY 

lost  +2  and  chlorine  lost  —2,  so  that  the  other  +  and  —  valences 
still  in  use  remained  equal  in  number,  and  equi-valence  was 
preserved. 

Balancing  an  Equation.  Suppose  we  wish  to  balance  the  equa- 
tion for  the  decomposition  of  chloric  acid  HC103.  We  ascertain, 
in  the  laboratory,  that  the  products  are  perchloric  acid  HC1O4, 
chlorine  dioxide  C102,  and  water. 

Skeleton:  HC103  ->  HC104  +  C102  +  H2O. 

Since  H+  and  0=  do  not  change  in  valence,  only  Cl  has  been 
affected.  On  the  left  side,  the  valences  are  Os  =  —  6,  H  =  +1,  Cl 
therefore  =  +5.*  On  the  right  side,  in  HC104,  the  total  valence 
of  oxygen  is  —Sand  of  hydrogen  +1.  That  of  Cl  is  therefore 
+7.  In  C102,  the  valence  of  02  is  —4,  and  that  of  Cl  there- 
fore +4.  Thus,  Cl  changes,  from  +5,  partly  to  +7  and  partly  to 
+4.  To  achieve  this,  arithmetically,  we  require  3C1  on  the  right 
(=  3  X  +5  =  +15),  giving  Cl  =  +7  and  2C1  =  2  X  +4  =  +8, 
or  a  total  of  +15  on  the  left.  Thus,  we  require  3HC10s : 

Balanced:  3HC103  =  HC104  +  2C102  +  H20. 

Balancing  Another  Equation.  In  the  reaction  for  preparing 
chlorine  (p.  157),  the  skeleton  is: 

Skeleton:     KMn04  +  HC1  ->  H20  +  KC1  +  MnCl2  +  Cl. 

Here,  in  KMn04,  the  valence  of  Mn  is  +7.  In  MnCl2  it  is  +2, 
a  loss  of  +5.  The  chlorine  also  changes  its  valence  from  —  1  to  0, 
a  loss  of  —  1.  Evidently,  so  that  the  changes  may  cancel  out,  for 
every  Mn  losing  +5,  5C1  must  lose  5  X  —1  and  be  liberated: 

Incomplete:      KMn04  +  HC1  -»  H20  +  KC1  +  MnCl2  +  5C1. 

Since  there  is  now,  altogether,  8C1  on  the  right,  8HC1  will  be  re- 
quired on  the  left.  The  8H  will  give  4H20: 

Balanced:       KMnO4  +  8HC1  ->  4H20  +  KC1  +  MnCl2  +  5C1. 
Molecular:  2KMnO4  +  16HC1  ->8H2O  +  2KC1  +  2MnCl2  +  5C12. 

For  another  method  of  balancing  this  equation,  see  p.  157. 
*  Write  these  (and  other  formula?)  thus:  H3+  Clft+O3~  (cf.  p.  276). 


OXIDATION   AND   REDUCTION  325 

Making  Equations.,  by  Using  the  Anhydrides.  —  To  balance 
the  equation  for  the  decomposition  of  chloric  acid,  we  first  write 
the  skeleton  equation: 

Skeleton:  HC103  -> HC104  +  C102  +  H20. 

Then  we  divide  the  acids  into  water  and  the  anhydrides  (p.  316). 
Analyzed:  H20,C1205  -*  H20,C1207  +  C102  +  H2O. 

We  now  perceive  that,  disregarding  the  water,  some  C1205  must 
lose  oxygen  to  give  2C102  +  0,  and  that  some  C12O5  must  gain  2O, 
becoming  C1207.  To  furnish  the  20,  clearly  2C1205  is  required, 
giving  4C1O2  +  20,  and  a  third  C12O5  gains  this  2O.  Thus,  alto- 
gether 3C1205  will  be  required: 

Balanced:       3H20,C1205  -»  H20,C1207  +  4C102  +  2H20 
or  6HC103  ->  2HC104  +  4C102  +  2H20. 

This  equation  is  then  divided  by  two  throughout. 

Making  Equations  by  Oxidation  of  Ions,  Using  Positive 
Electrical  Changes.  —  All  oxidation  reactions  involving  ionogens 
can  be  written  in  terms  of  ions.  Thus,  .the  oxidation  of  hydro- 
chloric acid  by  potassium  permanganate  can  be  so  written.  The 
potassium-ion  clearly  is  not  affected,  and  may  be  omitted.  The 
ions  concerned  are: 

MnOr  +  H+  +  OP  ->  H20  +  Mn++  +  Cl°. 

Cl°  with  no  charge  stands  for  free  chlorine.  Now  we  can  divide 
the  action  into  (1)  the  behavior  of  the  oxidizing  agent,  which 
is  general,  and  will  be  used  wherever  the  same  oxidizing  agent  is 
used;  (2)  the  fate  of  the  substance  being  oxidized,  which  again  is 
general,  because  other  oxidizing  agents  will  change  it  in  the  same 
way. 

MnO4~  +  8H+  ->  4H2O  +  Mn++  +  50 .  (1) 

In  words,  each  permanganate  ion,  with  a  free  acid  present  (oxi- 
dizing mixture),  will  give  water,  manganous-ion,  and  a  balance 
of  five  unit  positive  charges. 

50  +5Cr->5Cl°.  (2) 

100  +  5O2^  -+  5H20  +  50°2. 
100  +  5SO3=  +  5H20  -»  5SO4=  +  10H+ 


326  COLLEGE    CHEMISTRY 

These  three  equations  represent  the  oxidation  of  (2)  hydrochloric 
acid,  or  (21)  hydrogen  peroxide,  giving  free  oxygen,  or  (2n)  sul- 
phurous acid,  with  water  furnishing  the  oxygen,  and  leaving  the 
solution  strongly  acid  (=  5H2SO4).  Note  that  the  sums  of  the  + 
and  —  charges  on  opposite  sides  of  each  equation  are  equal. 
To  obtain  the  final  ionic  equation,  add  (1)  and  (2) : 

Mn04~  +  8H+  ->  4H20  +  Mn++  +  50 .          (1) 

5©  +  5C1-  -»  5C1°. (2) 

Mn04~  +  8H+  +  5C1~  ->  4H2O  +  Mn++  +  5C1°. 

Before  adding  (1)  and  (21)  and  (1)  and  (211),  the  first  equation  (1) 
must  be  doubled  throughout,  so  that  the  10©  may  cancel  out. 

Exercises.  —  1.   Assign  to  its  proper  class  (pp.  166,  258)  each 
of  the  actions  mentioned  in  this  chapter. 

2.  Knowing  that   potassium  fluosilicate  K2SiFe  is   insoluble, 
how  should  you  make  chloric  acid  (p.  314)? 

3.  Make  the  equation  for  the  interaction  of  chlorine  with 
calcium  hydroxide  in  hot  water  (p.  313) ,    How  should  you  make 
zinc  chlorate  from  zinc  hydroxide  Zn(OH)2? 

4.  How  should  you  make  pure  potassium  hypochlorite  from 
hypochlorous  acid  (p.  254)? 

5.  Explain,  in  terms  of  ionic  equilibrium,  why  dilute  hypo- 
chlorous  acid  can  be  obtained  by  adding  one-half  of  an  equivalent 
of  an  active  acid  (p.  309)  to  bleaching  powder,  and  distilling  the 
mixture. 

6.  On  what   circumstances  would  the  possibility  of  making 
barium  chlorate  by  action  of  chlorine  on  barium  hydroxide  depend 
(p.  313)?     Could  pure  barium  chlorate  be  obtained  easily  by  this 
means  (see  Table  of  Solubilities)? 

7.  Make  the  equations  for:    (a)  the  preparation  of  potassium 
bromate;    (6)  pure  aqueous  bromic  acid;    (c)  the  interaction  of 
iodine  with  aqueous  potassium  hydroxide  in  the  cold,  and  (d) 
when  heated. 

8.  Make  the  equations  for  the  interactions  of  chlorine  dioxide 
with  water,  and  with  aqueous  potassium  hydroxide. 

9.  Find  the  formulae  of  the  anhydrides  of  the  following  acids: 

HP03,    H2Se04,    H3As03,    H3As04,    H6S06. 


OXIDATION   AND   REDUCTION  327 

10.   Find  the  formulae  of  the  anhydrides  of  the  acids  from  the 
following  formulae  of  salts  : 


Na2Si03,    Na2HP04,    NaH2P03; 

11.  Classify  the  following  changes  as  oxidations  or  reductions. 
(a)  H2Cr207  -r*  H2CrO4  +  Cr03;  (6)  HMnO4  -»  Mn02;  (c)  I0-*!"; 
(d)  2H202-»2H20-fO2. 

12.  Using  positive  and  negative  valences,  determine  whether 
each  of  the  following  formulas  is  correct  or  incorrect:  Ca(Mn04)2, 
A1(C1O4)3,  Na2HI06. 

13.  Apply  each  of  the  three  methods  (pp.  322,  325)  of  writing 
equations  to  the  four  following  reactions:   (a)  chlorine-water  on 
bromine;    (6)   chlorine-  water  on  hydrogen  sulphide,  giving  free 
sulphur;    (c)  potassium  permanganate  and  free  acid  on  hydrogen 
sulphide,  giving  free  sulphur;    (d)  potassium  dichromate  and  free 
acid  (p.  224)  on  hydrogen  sulphide,  giving  the  chromic  salt  of  the 
acid  (Crm)  and  free  sulphur. 


CHAPTER  XXIV 


THE  ATMOSPHERE.     THE  HELIUM   FAMILY 

THE  pressure  which  is  exerted  by  the  air  upon  each  square  centi- 
meter of  the  earth's  surface  is  1033.6  g.,  or  a  little  over  one  kilo- 
gram. This  is  nearly  fifteen  pounds  to  the  square  inch. 

There  are  three  classes  of  components  in  the  air.  Those  of  the 
first  class,  oxygen,  nitrogen,  and  the  inert  gases  of  the  helium 
family,  are  present  in  almost  constant  quantities.  Those  of  the 
second  class  are  very  variable  in  quantity,  and  include  carbon 
dioxide,  water  vapor,  and  dust.  Those  of  the  third  class,  such  as 
the  sulphur  dioxide  in  city  air,  are  accidental. 

Components  which  are  Constant  in  Amount.  —  In  the 

determination  of  the  oxygen  in  air,  phosphorus  enclosed  in  iron 
gauze  (Fig.  90),  may  be  used.  The  oxygen  com- 
bines to  form  several  oxygen  acids  of  phosphorus. 
The  volume  of  gas  is  read  off  before  the  phosphorus  is 
introduced,  and  after  it  has  been  withdrawn. 

In  the  air  taken  from  mines,  from  mountain  tops, 
from  the  surface  of  the  sea,  and  from  inland  regions, 
the  percentages  of  oxygen  by  volume  are  found  to  be 
very  constant  (20.9  to  21.0). 

When  air,  from  which  the  oxygen  has  been  re- 
moved by  phosphorus,  or  by  passage  over  heated 
copper  or  iron,  is  led  slowly  through  a  heated  tube 
containing  magnesium,  the  nitrogen  unites  with  the 
metal  to  form  the  solid  magnesium  nitride  Mg3N2, 
and  only  about  10  c.c.  out  of  every  liter  remains 
uncombined.  This  residuum  is  argon,  mixed  with 

0.15  per  cent  of  its  volume  of  other  gases  belonging  to  the  helium 

family. 

The  Carbon  Dioxide.  —  Pure  country  air  contains  about  3 
parts  in  10,000  of  carbon  dioxide  C02.  In  city  air  there  are  from  6 

328 


Fia.  90. 


THE   ATMOSPHERE  329 

to  7  parts  in  the  same  volume,  while  in  the  air  of  audience-rooms 
the  proportion  may  rise  as  high  as  50  parts. 

The  sources  of  the  carbon  dioxide  in  the  air  are  numerous.  It 
comes  from  the  decay  of  vegetable  and  animal  matter,  in  which, 
chiefly  through  the  influence  of  minute  vegetable  organisms,  the 
carbon  is  oxidized  to  carbon  dioxide.  It  is  formed  also  by  the  com- 
bustion of  coal  and  wood,  but  the  thirteen  hundred  million  tons 
of  coal  burned  annually,  giving  three  times  that  weight  of  carbon 
dioxide,  would  add  only  one-six  hundredth  to  the  total  present  in 
the  air.  It  is  exhaled  by  animals,  being  produced  in  the  body  by 
oxidation  of  the  carbon  in  the  food  which  they  eat.  It  also  issues 
from  the  earth,  in  volcanic  as  well  as  in  other  neighborhoods.  The 
proportion  of  this  gas  in  the  air  would  naturally  increase  continu- 
ously, though  slowly,  as  the  result  of  these  processes,  were  it  not 
that  it  is  removed  just  as  continuously  by  the  action  of  growing 
plants  (see  p.  387),  which  use  it  as  food.  It  may  be  added,  also, 
that  carbon  dioxide,  being  a  soluble  gas,  is  contained  in  sea  water 
(dissolved  and  as  Ca(HCO3)2),  and  the  total  amount  in  the  ocean 
is  much  greater  than  that  in  the  air.  The  removal  by  plants  and 
by  sea  water  thus  keeps  the  proportion  in  the  air  fairly  constant. 

The  presence  of  carbon  dioxide  in  the  breath  may  be  shown 
very  quickly  by  blowing  through  a  tube  into  calcium  hydroxide 
solution  (limewater).  Calcium  carbonate  CaCOs  is  precipitated. 
We  draw  about  500  c.c.  of  air  into  our  lungs  at  each  breath,  or 
half  a  cubic  meter  per  hour.  In  the  lungs,  some  oxygen  is  re- 
moved, the  percentage  by  volume  falling  from  21  to  16,  and  we 
add  some  carbon  dioxide,  the  proportion  increasing  from  0.03  in 
country  air  to  about  4  per  cent.  A  candle  flame  is  extinguished 
by  exhaled  air,  because  the  maintenance  of  such  a  flame  requires 
at  least  18.5  per  cent  of  oxygen.  But  air  will  sustain  life  until  the 
proportion  has  fallen  to  about  10  per  cent. 

To  determine  the  proportion  of  carbon  dioxide,  a  measured  volume 
of  air  is  bubbled  slowly  through  a  measured  volume  of  a  solution 
of  barium  hydroxide  of  known  concentration.  Barium  carbonate  is 
precipitated:  Ba(OH)2  +  C02-*BaC03  j  +  H2O,  and  the  quantity 
of  barium  hydroxide  remaining  is  determined  by  titration  (p.  256). 

The  Water  Vapor.  —  The  proportion  of  water  vapor  is  con- 
stantly changing.  When  the  air  becomes  cool,  as  it  does  most  often 


330  COLLEGE    CHEMISTRY 

in  the  upper  layers,  the  vapor  condenses  to  droplets,  forming  fogs 
and  clouds.  When  the  condensation  continues,  the  drops  be- 
come larger  and  fall  as  rain.  On  the  other  hand,  when  the  weather 
is  warm,  water  from  the  soil,  and  from  rivers,  lakes,  and  oceans, 
passes  into  vapor  and  the  amount  in  the  air  increases. 

Humidity.  —  The  moisture  in  the  air  is  usually  denned  in 
terms  of  the  relative  humidity,  the  standard  being  the  quantity 
required  to  saturate  the  air.  The  open  air  is  never  actually 
saturated,  but,  when  a  portion  is  confined  in  a  vessel  over  water, 
it  soon  becomes  so.  The  humidity  is  then  100  per  cent.  If  the 
partial  pressure  of  water  vapor  present  is  only  half  as  great  as  the 
vapor  pressure  of  water  at  the  same  temperature,  the  humidity  is 
50  per  cent.  The  average  humidity  is  roughly  about  66  per  cent. 

At  18°  (64.4°  F.),  the  vapor  pressure  of  water  is  15.4  mm.  Thus 
air  saturated  with  moisture  at  18°  (100  per  cent  humidity)  would 
contain  15.4/760,  or  about  2  per  cent  by  volume  of  water  vapor. 
If  this  air  were  cooled  to  0°  (32°  F.),  a  temperature  at  which  the 
vapor  pressure  of  water  is  only  4.6  mm.,  the  air  could  retain  only 
4.6/760,  or  0.6  per  cent,  of  moisture.  The  difference,  amounting 
to  10.4  g.  (10.4  c.c.)  of  water  per  cubic  meter,  would  condense  as 
fog  or  rain. 

The  proportion  of  water  in  a  given  volume  of  air  may  be  meas- 
ured most  accurately  by  permitting  the  air  to  stream  slowly 
through  tubes  filled  with  calcium  chloride  or  phosphoric  anhydride. 
The  increase  in  weight  of  the  charged  tubes  represents  the  quantity 
of  moisture  abstracted  from  the  sample.  It  may  also  be  ascer- 
tained by  noting  the  temperature  to  which  air  has  to  be  cooled 
before  it  becomes  saturated  and  deposits  dew  (dew-point).  For 
example,  if  air  at  18°  has  to  be  cooled  to  11°  before  it  deposits  dew, 
it  contains  water  vapor  at  a  pressure  of  9.8  mm.  (Appendix  IV). 
If  saturated  at  18°,  it  would  have  contained  water  vapor  at  a 
partial  pressure  of  15.4  mm.  The  relative  humidity  was,  there- 
fore, 9.8/15.4,  or  63.6  per  cent. 

Ventilation.  —  On  a  moist  day,  we  speak  of  the  atmosphere 
as  "heavy"  or  " oppressive."  The  barometer,  however,  is  lower 
on  such  days,  and  the  pressure  below  the  average.  Moist  air 
must  be  lighter  than  dry  air,  because  in  moist  air  molecules  of 


THE   ATMOSPHERE  331 

relative  weight  18  (H20)  have  been  substituted  for  an  equal  num- 
ber of  molecules  of  oxygen  and  nitrogen  with  the  relative  weights 
32  and  28.  The  discomfort  is  due  to  a  different  cause. 

The  oxidation  of  digested  food  carried  by  the  blood  is  accom- 
panied by  liberation  of  heat,  yet  our  bodies  must  remain  at  98.6°  F. 
(37°  C.).  A  rise  of  a  few  tenths  of  a  degree  produces  discomfort. 
A  little  of  the  heat  is  lost  by  radiation  from  the  surface  of  the  body, 
but  the  real  adjustment  is  secured  by  evaporation  of  water  through 
the  skin.  The  vaporization  of  1  g.  of  water  (at  100°)  removes  heat 
amounting  to  540  calories  (603  cal.  at  37°  C.).  Evaporation  of  a 
single  ounce  (28f  g.)  of  water  will  therefore  lower  the  temperature 
of  96.5  kilograms  (168  Ibs.)  of  water  (or  flesh,  which  is  largely 
water)  by  more  than  two-tenths  of  a  degree  C.  (nearly  0.4°  F.). 

The  "oppressive"  feeling,  then,  is  due  to  the  fact  that  the  air 
is  too  nearly  saturated,  evaporation  is  being  hindered  (p.  90), 
and  heat  is  accumulating.  Hence,  the  relative  humidity  is  the 
measure  of  the  goodness  or  badness  of  the  air  of  a  room. 

In  winter,  cold  and  therefore  relatively  dry  air  is  brought  into 
the  house  and  heated.  This  makes  the  relative  humidity  very 
low,  evaporation  proceeds  too  fast,  and  discomfort  follows.  In 
summer,  however,  the  outside  air  is  often  already  nearly  satu- 
rated at  the  temperature  of  the  room.  Unless  there  is  a  rapid 
change  of  air  by  ventilation,  the  moisture  from  the  bodies  of  those 
in  the  room  increases  the  humidity,  and  discomfort  arises  from  a 
cause  opposite  to  the  one  which  produced  it  in  winter. 

It  should  be  noted,  also,  that  even  though  the  air  is  in  constant 
motion,  the  layer  of  air  next  our  skin  (even  the  exposed  parts)  is 
hindered  from  moving  by  friction.  There  is  a  stationary  layer 
close  to  the  surface,  which  quickly  reaches  the  temperature  of  the 
body  and  becomes  saturated  at  that  temperature.  The  water 
molecules  can  leave  this  layer,  and  make  room  for  others,  only  by 
diffusion,  which  is  a  deliberate  rather  than  a  speedy  process. 
Now,  an  electric  fan,  although  it  brings  no  fresh,  dryer  air  into  the 
room,  nevertheless  stirs  the  air  and  blows  away  the  moist,  saturated 
layer  next  the  skin.  It,  at  least,  makes  this  layer  much  thinner, 
and  reduces  greatly  the  distance  the  water  molecules  have  to  go 
by  mere  diffusion.* 

*  The  same  conception  applies  to  dissolving  a  salt.  A  stationary  layer  of 
saturated  solution  is  formed  on  the  surface,  and  the  molecules  of  the  salt  can 


332  COLLEGE    CHEMISTRY 

Formerly,  the  accumulation  of  carbon  dioxide  from  the  breath 
was  blamed  for  the  unhealthiness  of  unventilated  rooms.  The 
proportion  found  in  such  rooms,  however,  is  almost  never  sufficient 
to  do  any  harm.  Then,  it  was  imagined  that  traces  of  highly 
poisonous  compounds  were  exhaled  by  the  body.  No  one,  how- 
ever, has  yet  been  able  to  prove  that  such  poisons  exist. 

The  aims  of  ventilation  are,  therefore,  to  supply  fresh  outside 
air,  to  keep  it  in  motion,  and  to  maintain  a  humidity  that  is 
neither  too  low  nor  too  high. 

Dust  in  the  Air.  —  A  beam  of  sunlight,  crossing  a  dark  room, 
can  be  seen  by  the  light  reflected  from  the  particles  of  dust  in  the 
air.  Some  of  the  particles  are  inorganic,  and  consist  of  clay, 
limestone,  and  soot  from  ill-burned  fuel.  The  organic  du^t  may 
be  divided  into  two  kinds.  The  part  which  is  dead  includes  coal 
dust,  refuse  from  the  streets,  minute  shreds  of  cotton,  linen,  hay, 
etc.  The  living  dust  consists  of  pollen  grains,  spores  of  fungi 
and  molds,  bacteria,  and  similar  microscopic  organisms.  The 
presence  of  microscopic  germs  in  the  air  is  shown  by  the  fact  that 
when  food  has  been  exposed  to  the  air,  even  for  a  few  minutes, 
putrefaction  very  soon  sets  in.  Some  germs  also  produce  disease 
when  they  land  on  a  place  where  the  skin  has  been  damaged  by  a 
cut  or  a  burn.  After  infection,  antiseptic  treatment,  e.g.,  with 
hydrogen  peroxide,  destroys  the  organisms.  But  protection,  e.g., 
with  petrolatum  (p.  391),  until  a  new  skin  has  formed,  is  better. 

It  is  worth  noting  that  natural  soil  contains  about  100,000 
micro-organisms  per  c.c.,  good,  unfiltered  river  water  from  6000  to 
20,000  per  c.c.,  and  pure  air  only  4  or  5  per  liter. 

If  dust  were  absent  from  the  air,  there  would  be  no  clouds  or 
rain.  Aitken  has  shown  that  the  water  vapor  will  not  condense  to 
fog  in  air  that  has  been  freed  from  dust  by  nitration.  When  moist 
air  is  cooled,  the  dust  particles  act  as  nuclei,  round  which  the 
liquid  grows  at  the  expense  of  the  vapor.  In  the  absence  of  dust, 
the  cooling  would  produce  supersaturation,  which  would  be  slowly 

escape,  and  make  room  for  more,  only  by  diffusion.  In  liquids,  this  is  a  very 
slow  process.  By  shaking  the  solid  and  liquid,  however,  the  stationary  layer 
is  partly  washed  away.  It  is  made  thinner,  so  that  the  distance  the  mole- 
cules have  to  travel  by  diffusion  is  greatly  reduced,  and  the  whole  operation 
is  hastened. 


THE   ATMOSPHERE  333 

relieved  by  condensation  on  the  surfaces  of  houses,  plants,  animals, 
and  land.  Thus,  in  a  dustless  atmosphere  an  awning  or  umbrella 
would  afford  no  shelter. 

The  formation  of  fog  in  ordinary  air,  and  its  absence  in  filtered 
air —  e.g.,  air  drawn  through  a  wide  tube  packed  with  20-30  inches 
of  cotton  —  is  easily  shown  in  a 
darkened  room  (Fig.  91).  The 
flask  contains  some  water  to  satu- 
rate the  air.  When  suction  is 
applied,  by  the  mouth,  to  the  tube 
Sj  the  saturated  air  in  the  flask 
expands  and  is  cooled.*  With  ordi- 
nary air,  a  fog,  brilliantly  illumi- 
nated by  the  beam  of  light,  is 
instantly  produced.  Filtered  air 
(dustless)  gives  no  fog.  On  the 
other  hand,  a  whiff  of  smoke  from 

smoldering  paper,  when  admitted  to  the  flask,  causes  a  fog  (after 
cooling)  of  extraordinary  denseness. 

Composition  of  Air.  —  Air,  when  freed  from  carbon  dioxide 
and  water,  contains  by  volume  78.06  per  cent  of  nitrogen,  21.00 
per  cent  of  oxygen,  and  0.94  per  cent  of  argon.  When  only  the 
water  is  removed,  the  carbon  dioxide  averages  about  0.03  per  cent 
of  the  whole. 

To  use  an  illustration  of  Graham's,  if  we  imagined  the  air  to  be 
divided  by  magic  into  layers,  all  at  one  atmosphere  pressure,  and 
with  the  heavier  components  below,  we  should  have :  On  the  earth, 
five  inches  of  water;  above  that,  thirteen  feet  of  carbon  dioxide; 
above  that,  ninety  yards  of  argon;  above  that,  one  mile  of  oxygen; 
and  on  the  top  four  miles  of  nitrogen. 

Air  a  Mixture.  —  The  experiments,  in  which  the  oxygen  was 
removed  from  the  air  and  the  nitrogen  remained,  do  not  prove 
that  the  original  constituents  were  present  simply  in  mechanical 
mixture.  They  might  have  been  combined,  and  the  combustion 
of  phosphorus,  for  example,  might  have  represented  the  removal  of 
oxygen  from  combination  with  nitrogen  and  its  appropriation  by 
*  Compression  with  a  bicycle  pump  heats  air,  and  expansion  cools  it. 


334 


COLLEGE    CHEMISTRY 


the  phosphorus.     It  may  be  well,  therefore,  to  point  out  some 
reasons  which  lead  us  to  regard  the  air  as  a  mixture : 

1.  Each  of  the  substances  in  air  has  precisely  the  same  properties 
which  it  exhibits  when  free,  separate,  and  pure.     This  is  char- 
acteristic of  a  mixture.     Thus,  the  density  of  air  is  precisely  that 
which  we  find  by  calculation  from  the  known  proportions  and 
several  densities  of  the  components.     Again,  the  solubility  of  each 
gas  is  observed  to  be  the  same  as  if  the  same  amount  of  it  were 
present,  alone,  in  the  same  volume. 

2.  When  liquefied  air  is  allowed  to  evaporate  in  a  suitable 
apparatus,  the  nitrogen,  being  more  volatile,  can  be  separated 
completely  from  the  oxygen.     When  the  oxygen,  in  turn,  is  allowed 
to  evaporate,  the  carbon  dioxide  and  water  remain  as  solids, 

frozen  at  this  low  temperature. 

3.  Finally,  the  proportions  by  weight 
cannot  be  represented  by  a  chemical 
formula,  because  they  are  not  exact 
multiples  of  the  atomic  weights  by 
integral  numbers.  This  is  a  sure  proof 
that  it  is  not  a  chemical  aggregate. 

Liquefaction    of    Gases.  —  The 

earliest  experiments  of  this  kind  were 
made  by  Northmore  (1805),  who  lique- 
fied chlorine,  hydrogen  chloride,  and 
sulphur  dioxide.  In  1823  chlorine  was 
again  liquefied  by  Faraday.  During 
the  following  years  he  reduced  sulphur 
dioxide,  hydrogen  sulphide,  carbon  di- 
oxide, nitrous  oxide,  cyanogen,  and 
ammonia  to  the  liquid  condition.  He 
failed,  however  with  oxygen,  hydrogen, 
and  nitrogen.  In  1883  Wroblevski  and 
Olszevski  prepared  visible  amounts  of 
liquid  oxygen.  About  the  same  time  Dewar  devised  means  of  manu- 
facturing large  quantities  of  liquid  air  and  oxygen.  The  most  suc- 
cessful apparatus  for  use  on  a  small  scale  is  that  devised  by  Hampson. 
In  Hampson's  apparatus  (Fig.  92),  two  concentric  copper  pipes, 
about  130  meters  in  length,  are  coiled  closely  in  a  cylindrical  form, 


FIG.  92. 


THE   HELIUM   FAMILY  335 

with  non-conducting  covering  to  prevent  access  of  heat.  Air  at 
130-150  atmospheres  pressure  is  forced  through  the  inner  pipe. 
When  it  reaches  the  extremity  of  this  pipe,  it  suddenly  escapes  into 
a  closed  vessel.  This  expansion  lowers  its  temperature.  The  air 
can  now  escape  only  by  traveling  back  through  the  outer  pipe  to 
the  final  exit  near  the  top.  In  doing  so,  it  cools  the  highly  com- 
pressed air  in  the  inner  pipe.  This  cooler  air,  on  reaching  the 
closed  vessel,  expands  and  becomes  colder  than  ever,  and  in 
passing  backwards  lowers  the  temperature  of  the  air  in  the  inner 
pipe  still  further.  Finally,  the  air  in  this  pipe  liquefies,  and  drops 
of  liquid  air  are  expelled  into  the  closed  vessel.  They  are  allowed 
to  run  out  through  a  valve,  from  time  to  time,  as  they  accumulate. 

The  cooling  on  expansion  depends  upon  the  imper- 
fection (p.  78)  of  the  gas,  and  is  due  to  the  work 
done  in  overcoming  the  tendency  to  cohesion  of  its 
molecules.  Liquid  air  can  be  kept  in  Dewar  flasks 
(Fig.  93).  The  space  between  the  inner  and  outer 
flasks  is  evacuated,  so  that  there  is  no  gas  to  carry 
heat  to  the  liquid  air.  The  inner  surface  of  the  outer 
flask  is  often  silvered,  so  that  radiant  heat,  from  surrounding 
bodies,  may  be  reflected  and  not  absorbed. 

Liquid  Air.  —  Liquid  air  varies  in  composition,  as  the  nitrogen 
(b.-p.  —194°)  is  less  condensible  than  the  oxygen  (b.-p.  —181.4°). 
It  boils  at  about  — 190°,  and  contains  about  54  per  cent  of  oxygen 
by  weight,  while  air  contains  23.2  per  cent.  By  allowing  evapora- 
tion to  go  on,  a  liquid  containing  75  to  95  per  cent  of  oxygen  is 
easily  obtained  (cf.  p.  26).  The  gas  secured  by  the  evaporation 
of  the  residue  is  pumped  into  cylinders  and  sold  as  compressed 
oxygen.  It  contains  about  3  per  cent  of  argon,  and  is  a  con- 
venient source  of  this  element.  Cartridges  made  of  granular 
charcoal  and  cotton  waste,  when  saturated  with  liquid  air,  have 
been  used  as  an  explosive  in  mining. 

THE  HELIUM  FAMILY 

Argon  A.  —  Cavendish  (1785)  sought  for  other  gases  in  air  by 
adding  more  oxygen,  passing  an  electric  discharge  to  cause  this 
gas  to  combine  with  the  nitrogen,  and  absorbing  the  product 
(NO2)  in  potassium  hydroxide  solution.  He  found  that  only 


336  t        COLLEGE    CHEMISTRY 

about  0.8  per  cent  of  inactive  gas  remained.  Since  the  quantity 
was  so  small,  and  the  spectroscope,  by  which  the  gas  e^en  in  small 
amounts  would  have  been  recognized  to  Be  nertif,  was  riot  invented 
until  much  later,  he  did  not  pursue  the  subject. 

A  century  later,  Lord  Rayleigh  observed  that,  while  specimens 
of  oxygen  and  other  gases  made  purposely  from  various  sources 
always  had  the  same  density,  nitrogen  was  an  exception.  One 
liter  of  nitrogen  made  from  air,  and  supposed  to  be  pure,  weighed 
3.2572  g.  When  the  gas  was  manufactured  by  decomposition  of 
five  different  compounds,  such  as  urea  and  certain  oxides  of  nitro- 
gen, the  mean  weight  of  a  liter  of  this  nitrogen  was  only  1.2505  g. 
The  difference,  amounting  to  nearly  7  mgm.,  was  very  much  greater  ' 
than  the  experimental  error.  The  suspicion  naturally  arose  that 
some  heavier  gas  was  present  in  natural  nitrogen.  Soon  after 
(1894),  Rayleigh  repeated  Cavendish's  experiment,  and  obtained 
argon.  Working  in  cooperation  with  him,  Professor,  now  Sir 
William  Ramsay,  obtained  the  same  gas  by  removal  of  the  greatly 
preponderating  nitrogen  by  means  of  magnesium  (p.  328).  The 
new  gas  had  a  molecular  weight  of  about  40,  and  was  Aheref ojei- 
more  than  one-third  heavier  than  nitrogp*i^f»7f  0#  A*tfc^/M(  w 
The  exact  density  of  argon  is  39.88/When  liquefied  if  boils  at 

- 186.9°,  and  the  colorless  solid  melts  at  - 189.5°.  The  solubility 
of  the  gas  in  water  (4  volumes  in  100)  is  two  and  one-half  times 
that  of  nitrogen.  It  has  not  been  found  to  enter  into  any  sort  of 

chemical  combination,  and  was  named  argon  on  this  account  (Gk., 
inactive).  The  physical  properties  show  that  the  molecules  of  the 

gas,  like  those  of  mercury  (p.  Ill),  are  monatomic. 

Helium  He.  —  In  1868  Lockyer  first  detected  an  orange  line 
in  the  spectrum  of  the  sun's  prominences  which  was  not  given  by 
any  terrestrial  substance  then  known.  The  line  was  so  con- 
spicuous that  it  was  attributed  to  the  presence  of  a  new  chemical 
element,  which  was  named  helium  (Gk.,  the  sun).  Ramsay,  in 
searching  for  sources  of  argon,  examined  a  gas  which  Hillebrand 
had  obtained  from  uraninite,  an  ore  of  uranium.  He  was  sur- 
prised to  find  (1895)  that  it  contained  a  large  proportion  of  a  very 
light  gas,  the  spectrum  of  which  was  identical  with  that  of  solar 
helium.  The  same  gas  is  found  in  small  amount  in  the  atmosphere. 
Helium  does  not  exhibit  any  tendency  to  enter  into  combination. 


THE   HELIUM   FAMILY  337 

It  is  monatomic  and  its  density  shows  that  its  molecular  weight  is 
4.     When  liquefied  by  Onnes,  it  boiled  at  -268.5°  (4.5°  Abs.). 

Neon  Ne,  Krypton  Kr,  and  Xenon  Xe.  —  By  liquefying 
atmospheric  argon,  using  liquid  air  to  cool  it,  and  distilling  the 
liquid,  Ramsay  (1898)  found  that  it  contained  helium,  along  with 
three  new  gases.  These  together  constituted  one-six  hundredth 
part  of  the  whole.  The  gases  were  named  neon  (Gk.,  new), 
krypton  (Gk.,  hidden),  and  xenon  (Gk.,  stranger).  These  gases 
are  all  entirely  inactive  chemically,  and  are  all  monatomic.  Their 
molecular  weights  are:  Neon,  20.2;  krypton,  82.9;  xenon,  130.2 
Niton  Nt  (radium  emanation,  q.v.),  molecular  weight  222.4,  also 
belongs  to  this  family. 

Exercises.  —  1.  A  sample  of  moist  air,  confined  over  water  at 
15°  and  760  mm.,  occupies  15  c.c.  It  is  mixed  with  20  c.c.  of  hydro- 
gen, and  the  mixture  is  exploded,  and  suffers  a  contraction  of  9.5 
c.c.  What  would  be  the  volume  of  the  oxygen  it£oniained  if 
measured  dry  at  0°  and  760  mm.?  ^  *  7  &<***&& 

2.  Calculate,  from  the  data  on  p.  333  and  the  densities,  the 
percentage  by  weight  of  the  three  principal  components  of  air. 

3.  Of  the  proofs  that  air  is  a  mixture  (p.  333),  which  show  that  '   j . 
no  part  of  the  components  is  combined,  and  which  that  the  com— — ^~" 
ponents  are  not  wholly  combined? 

4.  What  is  the  relation  between  heavier  clothing  and  the 
stationary  layer  of  air  next  the  skin? 

5.  From  the  data  given  on  p.  330,  calculate  the  weight  of  water 
vapor  in  1  cubic  meter  of  air  saturated  at  18°  and  at  0°,  respectively. 


CHAPTER  XXV 
NITROGEN  AND  AMMONIA 

NITROGEN  was  recognized  to  be  a  distinct  substance  by  Ruther- 
ford (1772),  Professor  of  Botany  in  the  University  of  Edinburgh, 
who  named  it  mephitic  air.  Scheele  showed  that  it  was  present 
in  the  atmosphere.  Lavoisier  recognized  it  to  be  an  element,  and 
named  it  azote  (Gk.,  without  life)  because  it  did  not  support  life. 
The  English  name  records  the  fact  that  it  is  an  important  con- 
stituent of  niter  KNOs. 

The  Chemical  Relations  of  the  Element  Nitrogen.  —  In 

compounds  with  hydrogen  and  the  metals  nitrogen  is  trivalent, 
while  in  those  containing  oxygen  and  other  negative  elements,  it 
is  frequently  quinquivalent.  It  is  a  non-metal,  for  its  oxides  are 
acidic  (p.  94)  .  Many  of  the  compounds  of  nitrogen  are  extremely 
active  and  interesting.  Those  of  them  which  we  have  to  discuss 
in  inorganic  chemistry  are  ammonia  NHs  and  nitric  acid  HNOs, 
and  several  related  substances. 

Occurrence.  —  Free  nitrogen  is  present  in  the  air.  The 
nitrates  of  potassium  and  sodium  are  found  in  Bengal  and  Chile, 
respectively.  Natural  manures,  such  as  guano,  contain  large 
quantities  of  nitrogen  compounds,  and  owe  their  value  as  fertilizers 
to  this  fact.  Nitrogen  is  a  constituent  of  the  proteins  (about  15 
per  cent  nitrogen)  of  vegetable  and  animal  matter. 

Preparation.  —  Nitrogen  containing  about  one  per  cent  of 
argon  is  obtained  by  burning  phosphorus  in  air,  or  by  passing  air 
over  heated  copper:  2Cu  +  02  —  >  2CuO.  For  commercial  pur- 
poses, it  is  obtained  by  evaporation  of  liquid  air. 

Pure  nitrogen  is  prepared  by  heating  ammonium  nitrite: 


338 


NITROGEN   AND   AMMONIA  339 

In  practice,  since  ammonium  nitrite  is  unstable  and  cannot  be 
kept  as  such,  strong  solutions  of  ammonium  chloride  and  sodium 
nitrite  are  mixed,  a  double  decomposition  results  in  the  formation 
of  ammonium  nitrite,  NH^Cl  +  NaN02  ^±  NI^NOa  +  Nad,  and 
this  breaks  up  when  heat  is  applied,  giving  nitrogen. 

We  may  also  prepare  nitrogen  by  the  oxidation  of  ammonia 
NH3,  passing  the  latter  over  heated  cupric  oxide  (see  p.  343),  or 
by  the  reduction  of  nitric  oxide  NO,  passing  this  gas  over  heated 
copper. 

Physical  Properties.  —  Nitrogen  is  a  colorless,  tasteless, 
odorless  gas,  as  we  should  expect  from  the  fact  that  air  possesses 
these  properties.  It  forms  a  colorless  liquid,  boiling  at  —194°, 
and  a  white  solid  (m.-p.  —214°).  The  solubility  in  water  (1.6 
vols.  in  100)  is  less  than  that  of  oxygen.  The  density  of  the  gas 
shows  the  formula  of  free  nitrogen  to  be  N2. 

Chemical  Properties.  —  Nitrogen  unites  with  few  elements 
directly.  At  ordinary  temperatures  it  is  almost  absolutely  in- 
different. When  passed  over  heated  lithium,  calcium,  magnesium, 
or  boron,  it  forms  nitrides,  in  which  it  is  trivalent.  These  have 
the  formulae  Li3N,  Ca3N2,  Mg3N2,  and  BN,  respectively.  Thus, 
when  magnesium  is  burned  in  the  air,  the  white  mass  which  is 
formed  contains  magnesium  nitride,  along  with  much  of  the 
oxide.  When  the  ash  is  moistened  with  water  in  a  covered  vessel, 
ammonia  can  be  smelt  and  can  be  detected  with  moist  litmus 
paper.  The  nitride  is  hydrolyzed: 

Mg3N2  +  6H20  ->  3Mg(OH)2  +  2NH3t . 

Nitrogen  combines  with  difficulty  with  hydrogen  to  form  am- 
monia NH3  and  with  oxygen  to  form  nitric  oxide.  The  actions 
will  be  discussed  under  the  compounds  themselves. 

One  case  of  direct  union  of  nitrogen  is  of  economic  importance. 
The  supply  required  by  most  plants  is  obtained  from  nitrogen 
compounds  contained  in  fertilizers,  or  equivalent  substances 
already  present  in  the  soil.  With  the  leguminosce  (peas,  beans, 
clover,  etc.),  however,  are  found  associated  certain  bacteria, 
which  flourish  in  nodules  upon  their  roots.  These  bacteria  have 
the  power  of  taking  free  nitrogen  from  the  air,  which  penetrates 


340  COLLEGE    CHEMISTRY 

the  soil,  and  producing  proteins.  The  nodules  often  contain 
over  five  per  cent  of  combined  nitrogen.  The  proteins,  by  the 
action  of  nitrifying  bacteria,  give  nitric  acid  which,  with  bases  in 
the  soil,  gives  nitrates.  These  are  soluble,  and  are  absorbed 
through  the  roots,  furnishing  the  nitrogen  needed  by  plants  to 
enable  them  to  construct  the  proteins  they  require. 

Compounds  of  Nitrogen  and  Hydrogen.  — The  commonest 
and  longest  known  of  these  substances  is  ammonia  NH3,  which 
was  first  described  by  Priestley  (1774)  and  named  "  alkaline  air." 
Curtius  discovered  hydrazine  N2H4  in  1889,  and  hydrazoic  acid 
HN3  in  1890.  Hydroxylamine  HONH2,  discovered  by  Lossen  in 
1865,  is  similar  to  ammonia  in  chemical  behavior. 

AMMONIA  NH3 

Ammonia  is  of  interest,  commercially,  because  large  amounts 
of  liquefied  ammonia  are  used, in  refrigeration,  because  much  is 
employed  in  the  manufacture  of  carbonate  of  soda,  and  because  its 
compounds  are  used  as  fertilizers. 

Manufacture.  —  Ammonia  is  formed  when  proteins  are  heated 
in  the  absence  of  air.  Thus,  it  was  formerly  obtained  by  the 
distillation  of  hoofs,  hides,  and  horns,  and  the  solution  in  water 
was  called  "  spirit  of  hartshorn."  Coal  contains  about  1  per  cent 
of  combined  nitrogen,  derived  from  the  proteins  of  the  original 
plants.  Hence,  when  coal  is  distilled  in  the  manufacture  of  coal 
gas  or,  on  a  much  larger  scale,  for  the  making  of  coke,  much  am- 
monia can  be  secured  by  washing  with  water  the  gases  which  are 
given  off.  The  solution  is  separated  from  the  tar,  lime  is  added 
to  combine  with  acids,  and  the  ammonia  gas  is  driven  out  by 
heating  and  passed  into  sulphuric  acid  or  hydrochloric  acid.  It 
gives  ammonium  sulphate  or  chloride  (see  below). 

In  Germany  80  per  cent  (1910)  of  the  coke  is  made  in  "by- 
product" coke  ovens,  in  which  the  ammonia  and  other  by-products 
are  collected  and  utilized;  in  the  United  States  83  per  cent  of  the 
coke  is  made  in  " beehive"  ovens,  in  which  the  vapors  are  simply 
burned.  Ammonium  sulphate  is  a  valuable  fertilizer  and  in 
1911,  in  the  United  States,  ammonia  capable  of  yielding  400,000 


NITROGEN   AND   AMMONIA  341 

tons  of  ammonium  sulphate  worth  24  million  dollars  was  burned 
by  the  cokemakers. 

The  distillation  of  coal  is  the  chief  source  of  commercial  am- 
monia. In  Scotland,  however,  oil-bearing  shale  is  distilled  to 
obtain  petroleum,  and  much  ammonia,  liberated  at  the  same 
time,  is  collected.  Formerly  it  was  allowed  to  escape  but,  in  the 
absence  of  a  protective  tariff,  the  competition  of  American  and 
Russian  petroleum  compelled  economy.  Now,  the  profit  on  the 
ammonium  sulphate  pays  the  whole  cost  of  mining  and  distilling 
the  shale. 

Synthetic  Ammonia.  —  The  Badische  Company  is  now 
manufacturing  ammonia  on  a  large  scale,  for  the  preparation  of 
explosives,  by  the  direct  union  of  nitrogen  and  hydrogen. 

N2  +  3H2  <F±  2NH3  +  2  X  12,200  cal. 

No  union  occurs  at  low  temperatures  and,  on  the  other  hand,  the 
action  is  reversible  and  exothermal,  so  that  at  700°  ammonia  is 
decomposed  almost  completely  (Van't  HofFs  law,  p.  188).  It  is 
necessary,  therefore,  to  use  a  lower  temperature  and  a  contact 
agent  —  such  as  specially  prepared  iron  —  to  hasten  the  action. 
Then,  too,  the  reaction  is  accompanied  by  a  diminution  in  vol- 
ume (4  vols.  — >  2  vols.),  and  is  therefore  assisted  by  using  the 
gases  under  a  pressure  of  185-200  atmospheres  (Le  Chatelier's 
law,  p.  190).  At  500°,  with  these  conditions,  about  8  per  cent 
of  the  gases  combine.  The  ammonia  is  dissolved  out  with  water, 
and  the  uncombined  gases  are  sent  through  the  process  again. 
The  required  hydrogen  may  be  obtained  by  one  of  the  com- 
mercial processes  (p.  56),  and  the  nitrogen  from  liquid  air. 

Preparation  in  the  Laboratory.  —  1.  A  mixture  of  slaked 
lime  and  some  salt  of  ammonium,  such  as  ammonium  chloride, 
either  with  or  without  water,  is  heated  in  a  flask  or  retort  pro- 
vided with  a  delivery  tube: 

Ca(OH)2  +  2NH4C1  fc?  CaCl2  +  2NH40H  <±  2NH3  +  2H2O. 

The  ammonium  hydroxide,  formed  by  the  double  decomposition, 
immediately  decomposes. 


342 


COLLEGE    CHEMISTRY 


2.  Warming  the  aqueous  solution  gives  a  steady  stream  of  the 
gas.  Since  the  gas  is  very  soluble  in  water,  it  is  collected  over 
mercury  or  in  an  inverted  jar  by  downward  displacement  of  air. 
In  both  methods  of  preparation,  it  is  dried  with  quicklime  (p.  475). 

Physical  Properties.  —  Ammonia  is  a  colorless  gas  with  a 
pungent,  characteristic  odor  familiar  in  smelling-salts.  The 
G.M.V.  of  the  gas  weighs  17.26  g.,  so  that  the  density  is  little 
more  than  half  that  of  air  (cf.  p.  101).  When  liquefied  it  boils 
at  —33°  and  the  solid  is  white  and  crystalline  (m.-p.  —77°). 
One  volume  of  water  dissolves  1300  volumes  of  the  gas  at  0°, 
and  783  volumes  at  16°.  The  35  per  cent  solution,  sold  as  "con- 
centrated ammonia,"  has  a  sp.  gr.  0.881.  The  whole  of  the  dis- 
solved gas  may  be  removed  by  boiling  (cf.  p.  145). 

Liquefied  ammonia  is  used  in  refrigeration.  In  evaporating  at 
—  33°  it  absorbs  330  cal.  per  gram.  Water  alone  has  a  greater 

,, heat  of  vaporization.  The  large  amount 

of  heat  is,  in  both  cases,  required  be- 
cause of  the  relatively  large  volume  of 
the  vapor  (due  to  low  molecular  weight) 
and  to  the  fact  that  both  liquids  are 
associated  (p.  206),  and  the  complex 
molecules  (NHa)2  and  (NH3)3  have  to 
be  decomposed.  To  freeze  1  gram  of 
water  at  0°,  79  cal.  have  to  be  re- 
moved. Thus  1  g.  of  liquid  ammonia 
will  comvert  4  g.  of  water  into  ice. 
Fig.  94  shows  one  arrangement  diagram- 
matically.  The  ammonia  gas,  obtained 
from  a  cylinder  of  liquid  ammonia,  is 
driven  by  the  pump  F  along  the  tube 
E  and  is  liquefied  in  the  tube  coiled  in 
the  tank  A  B.  Cold  water  circulating 
through  AB  removes  the  heat  produced 
by  the  compression  and  liquefaction  of  the  gas.  The  liquid 
ammonia  is  allowed  to  drip  through  the  stopcock  G  into  the  lower 
coil,  and  there  it  evaporates.  In  doing  so,  it  takes  heat  from  a 
30  per  cent  solution  of  calcium  chloride  in  water.  This  cooled 
brine  leaves  the  tank  at  D,  circulates  through  another  tank,  in 


FIG.  94. 


NITROGEN   AND   AMMONIA  343 

which  water-filled  ice  molds  are  suspended,  and  returns  to  C. 
When  used  for  cooling  storage-rooms  for  meat,  the  brine  circulates 
through  pipes  in  the  same  way.  The  machine  is  constructed  of 
iron,  because  copper  and  brass  are  corroded  by  ammonia. 

Chemical  Properties.  —  Ammonia,  as  we  have  seen,  is  not 
stable,  and  decomposes 'almost  completely  at  700°.  A  discharge 
of  sparks  from  an  induction  coil  (temperature  about 
2000°)  has  the  same  effect,  so  that  a  sample  of  the 
gas,  confined  over  mercury  in  a  closed  tube  (Fig.  95), 
may  be  shown  to  double  in  volume.  Every  two 
molecules  give  four: 


That,  even  at  this  temperature,  the  action,  being 
reversible,  is  still  incomplete,  can  be  shown  by 
introducing  a  few  drops  of  dilute  sulphuric  acid. 
The  trace  of  ammonia  remaining  combines  with  this 
acid,  forming  (NH4)2SO4  in  solution.  If  the  dis- 
charge is  continued,  further  traces  of  ammonia  are 
formed  and  absorbed,  until,  finally,  the  whole  gas 
disappears. 

Ammonia  reduces  many  oxides,  when  the  latter 
are  heated  and  the  gas  is  led  over  them: 

3CuO  +  2NH3  ->  3Cu  +  3H2O  +  N2. 


Ammonia  burns  in  pure  oxygen  (not  in  air)  to  give 
steam  and  nitrogen. 

Chlorine  and  bromine  (vapor)  combine  with  the  hydrogen  and 
liberate  nitrogen: 

2NH3  +  3C12  ->  N2  +  6HC1. 

When  metals  capable  of  uniting  with  nitrogen  (p.  339)  are 
heated  in  a  stream  of  ammonia  gas,  hydrogen  is  displaced.  Mag- 
nesium gives  magnesium  nitride: 

2NH3  +  3Mg  ->  Mg3N2  +  3H2. 

Sodium  and  potassium,  however,  give  amides  (compounds  con- 
taining the  group  NH2),  such  as  sodamide  NaNH2: 
2NH3  +  2Na  ->  2NaNH2  +  H2. 


344  COLLEGE    CHEMISTRY 

The  most  striking  property  of  ammonia  is  that  it  combines  with 
acids,  giving  ammonium  salts: 

NH3  (gas)  +  HC1  (gas)     -»  NI^Cl  (solid). 
2NH3  (gas)  +  H2S04  (liq.)  -»  (NH4)2S04  (solid). 

It  combines  also  with  water  at  or  below  —  79.3°  to  give  ammonium 
hydroxide,  a  white  solid: 

NH3  +  H2O 


As  the  solid  dissociates  above  —79.3°,  a  solution  of  the  substance, 
which  is  contained  in  the  aqueous  solution  of  ammonia,  is  the 
only  available  form  of  ammonium  hydroxide.  In  solution,  it  is  a 
weak  base. 

Ammonium  oxide  (NH^O,  a  solid,  can  also  be  formed  below 
-78.6°. 

Ammonium  Compounds.  —  Since  NH4  plays  the  part  of  a 
metallic  element,  entering  into  the  composition  of  a  base  and  of 
a  series  of  salts,  it  is  named  ammonium.  As  this  radical  forms  a 
univalent,  positive  ion  NHf*"  and  gives  a  distinctly  alkaline  base, 
it  is  classed  with  the  metallic  elements  of  the  alkalies  (q.v.)  . 

Ammonium  Hydroxide.  —  Although  less  completely  ionized 
than  potassium  hydroxide,  ammonium  hydroxide  affects  litmus 
easily.  In  a  normal  solution,  0.4  per  cent  of  the  ammonia  is  in 
the  form  of  ammonium-ion  NH4+.  When  an  acid  is  added  to  the 
solution,  the  equally  small  amount  of  hydroxide-ion  which  exists 
in  it  is  removed  and  the  various  equilibria  are  displaced  forward. 
The  final  result  is  the  same  as  with  any  other  base: 


NH3(gas)  24  „ 

+  H+J- 


Probably  only  a  small  proportion  of  the  gas  is  actually  com- 
bined at  any  one  time,  the  greater  part  being  simply  dissolved. 

The  solution  is  sold  as  household  ammonia,  and  is  used,  in 
washing  and  cleaning,  to  soften  the  water. 

Salts  of  Ammonium.  —  When  strongly  heated,  all  am- 
monium salts  are  decomposed  and  many,  but  not  all,  give  am- 


NITROGEN   AND   AMMONIA  '345 

™ 

monia  and  the  acid.  When  the  latter  is  volatile,  the  whole 
material  of  the  salt  is  thus  converted  into  gas.  The  acid  and  the 
ammonia  reunite  to  form  the  solid  salt  when  the  vapor  reaches  a 
cool  part  of  the  tube  (sublimation,  p.  199) : 

NH4C1  (solid)  <=>  NH4C1  (gas)  <=»  HC1  +  NH3. 

The  test  for  ammonium  salts  is  to  warm  them,  dry  or  in  solution, 
with  a  base,  when  the  odor  of  ammonia  becomes  noticeable. 

(NH4)2SO4 1=?  SO4=  +  2NH4+ 
2KOH  fc>  2K+  +  20BT 

When  the  solution  is  used,  it  is  the  tendency  of  the  NHt+  and 
OH~  to  unite  to  form  the  slightly  ionized,  molecular  hydroxide 
that  sets  the  other  equilibria  in  motion. 

In  ammonium  salts,  the  nitrogen  is  quinquivalent. 

Hydrazine  N2H4:.  —  By  reduction  of  a  compound  of  nitric  oxide 
and  potassium  sulphite  by  means  of  sodium  amalgam,*  a  solution 
of  hydrazine  hydrate  is;  obtained: 

K2S03,2NO  +  3H2  ->  N2H4,H20  +  K2S04. 

When  the  hydrate  is  distilled  with  barium  oxide,  under  reduced 
pressure,  hydrazine  is  liberated: 

N2H4,H20  +  BaO  ->  N2Hit  +  Ba(OH)2. 

Hydrazine  hydrate  freezes  at  about  —40°  (b.-p.  118.5°).  Its 
aqueous  solution  is  alkaline,  and  salts  are  formed  by  neutraliza- 
tion. 

Hydrasoic  Acid  HN3.  —  When  nitrous  oxide  (q.v.)  "is  led  over 
sodamide  at  200°,  water  is  liberated  and  sodium  hydrazoate  re- 
mains behind: 

NH2Na  +  N2O  ->  NaN3  +  H20^ , 

A  dilute  solution  of  the  free  acid  is  best  obtained  by  distilling  the 
lead  salt  with  dilute  sulphuric  acid. 

The  pure  acid  (b.-p.  37°)  is  violently  explosive,  resolving  itself 
into  nitrogen  and  hydrogen  with  liberation  of  much  heat: 
2HN3,Aq  ->  H2  +  3N2  +  Aq  +  2  X  61,600  cal. 

*  The  sodium  dissolved  in  the  mercury  interacts  with  the  water,  ^giving 
hydrogen  (see  Active  state  of  hydrogen). 


346  COLLEGE    CHEMISTRY 

Halogen  Compounds  of  Nitrogen.  —  When  ammonium 
chloride  solution  is  treated  with  excess  of  chlorine,  drops  of  an 
oily  liquid,  nitrogen  trichloride,  are  formed:  3C12  +  NI^Cl  — •» 
NCls  +  4HC1.  It  is  extremely  explosive,  resolving  itself  into  its 
constituents  with  liberation  of  much  heat. 

When  a  solution  of  iodine  in  potassium  iodide  solution  (p.  200) 
is  added  to  aqueous  ammonia,  a  brown  precipitate  is  formed. 
This  seems  to  have  the  composition  NH3,NIa,  and  is  named 
nitrogen  iodide.  It  may  be  handled  while  wet.  When  dry,  if 
touched  with  a  feather,  it  decomposes  into  its  constituents  with 
violent  explosion. 

' r~~^~-^ 

Exercises.  —  1.  When  moist  air  is  used  as  a  source  of  nitrogen, 
what  advantage  is  there  in  using  copper  rather  than  the  less 
expensive  metal  iron,  for  removing  the  oxygen  (p.  60)? 

2.  How  many  grams  of  water  at  0°  could  be  frozen  (p.  85)  by 
the  removal  of  the  heat  required  to  evaporate  50  g.  of  liquid 
ammonia?  £  0  }  ^  -  ^  p^  £>^ 

How  many  grams  of  ammonia  are  containedm  1  1.  of  "  con- 
centrated ammonia"  (p.  342)? 

4.  What    are    the    ions    of    hydrazine    hydrate?     Formulate 
(p.  254)  the  neutralization  of  this  base  with  sulphuric  acid. 

5.  What  is  the  object  attained  by  distilling;  under  reduced 
pressure  in  making  hydrazine  (p.  345)?         ()  *  \*  /  f  f   /(} 

6.  Classify  (pp.  166,  258),  (a)  the  interaction  of  a  nitride  with 
water  (p.  343)  and  (6)  of  chlorine  and  ammonium  chloride  (p.  343), 

^^    (c)  the  results  of  heating  ammonium  nitrite  (p.  338)  and  (d)  am- 

^^monium  chloride  (p.  345). 
*V        7.  Why  does  not  ammonia  burn  in  air  (p.  343)? 

>3        8.   What    substances    are    present    in    ammonium    hydroxide 
solution?     When  the  liquid  is  heated,  what  happens  to  each? 

|j     Formulate  the  system. 


CHAPTER  XXVI 
OXIDES  AND   OXYGEN  ACIDS   OF  NITROGEN 

THE  names  and  formulae  of  the  oxides  and  oxygen  acids  of 
nitrogen  are  as  follows: 

Nitrous  oxide  N2O  < Hyponitrous  acid  H2N202 

Nitric  oxide  NO 

Nitrous  anhydride  N203          < >    Nitrous  acid  HNO2 

Nitrogen  tetroxide  N2C>4  and  NO2 

Nitric  anhydride  N20s  < >    Nitric  acid  HNOs. 

All  the  oxides  are  endo thermal  compounds  (p.  174),  yet,  with  the 
exception  of  the  third  and  the  last,  they  are  all  relatively  stable. 
The  acids,  when  deprived  of  the  elements  of  water,  yield  the  oxides 
opposite  which  they  stand  (p.  281,  footnote).  Conversely,  ex- 
cepting in  the  case  of  nitrous  oxide,  the  anhydrides  with  water  give 
the  acids.  All  of  these  substances  are  made  directly  or  indirectly 
from  nitric  acid  —  nitric  anhydride  by  removal  of  water,  the 
others  by  reduction.  We  turn,  therefore,  first,  to  nitric  acid  and 
its.  properties.  This  acid  is  made  from  Chile  saltpeter  (next  sec- 
tion) and  also  by  fixation  of  atmospheric  nitrogen  (see  p.  352). 

NITRIC  ACID  HNOs 

Sources.  —  Sodium  nitrate,  or  Chile  saltpeter  (caliche)  is 
found  in  a  desert  region  near  the  boundary  of  Chile  and  Peru. 
The  deposit  is  about  5  feet  thick,  2  miles  wide,  220  miles  long,  and 
contains  20  to  60  per  cent  of  the  salt.  Purification  is  effected  by 
recrystallization.  Potassium  nitrate,  or  Bengal  saltpeter,  is  found 
in  the  soil  in  the  neighborhood  of  cities  in  India,  Persia,  and  other 
oriental  countries.  It  arises  from  the  oxidation  of  animal  refuse 
to  nitric  acid,  through  the  mediation  of  nitrifying  bacteria.  The 
potash  and  lime  in  the  soil,  along  with  the  nitric  acid,  give  nitrates 
of  potassium  and  calcium.  The  aqueous  extract  of  this  soil 
is  treated  with  wood  ashes,  which  contain  potash  K2COs.  It  is 

347 


348  COLLEGE    CHEMISTRY 

poured  off  from  the  calcium  carbonate  thus  precipitated,  and  is 
finally  evaporated.  In  guano  (excreta  of  sea  birds),  used  as  a 
fertilizer,  the  nitrogen  compounds  have  often  been  converted 
largely  into  nitrates  in  the  same  way. 

Manufacture.  —  When  any  nitrate  is  treated  with  any  acid, 
nitric  acid  is  formed  by  a  reversible  double  decomposition.  As 
sodium  nitrate  is  the  cheapest  salt  of  nitric  acid,  it  is  always  em- 
ployed. For  the  same  reason  and,  above  all,  because  of  its  relative 
in  volatility,  sulphuric  acid  is  used  to  displace  it: 

NaN03  +  H2S04  t*  NaHS04  +  HN03T. 

The  nitric  acid  is  rather  volatile  (b.-p.  86°),  while  sulphuric  acid 
(b.-p.  330°)  is  much  less  so,  and  the  two  salts  are  not  volatile  at  all. 
The  materials  are  heated  in  cast-iron  stills,  and  the  nitric  acid 
vapor  is  condensed  in  glass  pipes  surrounded  by  water.  Thus  the 
interaction  proceeds  to  completion  very  easily  (cf.  p.  142;  see  also 
p.  185). 

Physical  Properties.  —  Nitric  acid  is  a  colorless,  mobile  liquid 
(sp.  gr.  1.52)  boiling  at  86°,  and  freezing  to  a  solid  (m.-p.  —47°). 
It  fumes  strongly  when  its  vapor  issues  into  moist  air  (cf.  p.  144). 
An  aqueous  solution  containing  68  per  cent  of  the  acid  boils  at 
120.5°,  while  the  pure  acid,  pure  water,  and  all  other  mixtures, 
boil  at  lower  temperatures.  This  68  per  cent  nitric  acid  of  constant 
boiling-point  (p.  145)  forms  the  "concentrated  nitric  acid"  of 
commerce  (sp.  gr.  1.41). 

Chemical  Properties.  —  1.  Like  chloric  acid  (p.  314),  and 
other  oxygen  acids  of  the  halogens,  nitric  acid  is  most  stable  when 
mixed  with  water.  The  pure  (100  per  cent)  acid  decomposes  while 
being  distilled: 

4HNO3  ->  4NO2  +  2H2O  +  O2, 

yet  not  with  explosive  violence  like  chloric  acid.  The  distillate 
is  colored  brown  by  dissolved  nitrogen  tetroxide  NO2  ("  fuming  " 
nitric  acid).  Repeated  distillation  finally  leaves  68  per  cent  of  the 
acid,  mixed  with  32  per  cent  of  water  formed  by  the  above  decom- 
position. The  acid  of  constant  boiling-point  is,  therefore,  reached, 


OXIDES   AND   OXYGEN  ACIDS   OF  NITROGEN  349 

as  usual,  from  more  concentrated  as  well  as  from  less  concentrated 
specimens. 

2.  Nitric  acid,  when  dissolved  in  water,  is  highly  ionized,  and 
is  therefore  active  as  an  acid.    By  interaction  with  hydroxides  and 
oxides  it  forms  nitrates. 

3.  When  pure  nitric  acid  (b.-p.  86°)  is  poured  upon  phosphoric 
anhydride,  the  latter  combines  with  the  elements  of  water,  and  dis- 
tillation gives  nitric  anhydride  :   2HN03+  P2O5  ->  N205  T  +  2HPO3. 
The  anhydride  is  a  white  solid  melting  at  30°  and  boiling  at 
45°.     It  unites  vigorously  with  water  to  form  nitric  acid.     It 
decomposes  spontaneously  into  nitrogen  tetroxide  and  oxygen: 
2N205->4N02  +  O2. 

4.  Like  the  unstable  oxygen  acids  of  the  halogens,  nitric  acid  is 
an  oxidizing  agent  even  when  diluted  with  water.     The  multiplicity 
of  the  products  into  which  it  may  be  decomposed  by  reduction, 
however,  renders  separate  treatment  of  this  property  necessary 
(see  p.  354). 

5.  Nitric  acid  interacts  energetically  with  many  compounds  of 
carbon  to  give  nitro-derivatives.     Thus,  when  heated  with  phenol 
CeH^OH)    (carbolic   acid)   it   gives   picric   acid    (trinitrophenol) 
C6H2(N02)3(OH),  which  crystallizes  in  yellow  needles  in  the  mix- 
ture.    This  is  a  yellow  dye,  used  also  as  an  explosive. 

C6H5(OH)  +  3HON02  -»  C6H2(OH)(N02)3  +  3H20. 


When  heated  with  toluene  CeHsCHs,  it  gives  trinitrotoluene: 
CH3C6H5  +  3HON02  -»  CH3'c6H2(N02)3  +  3H20. 

This  substance  (T.N.T.)  is  used  for  filling  "high  explosive"  shells, 
because  it  can  be  melted  (m.-p.  81.5°)  and  poured  in,  making  the 
filling  easy,  safe,  rapid,  and  complete.  It  is  not  easily  exploded  by 
shocks  during  transportation,  but  it  explodes  instantaneously  and 
completely  with  a  detonator.  The  following  equation  shows, 
roughly,  the  decomposition,  and  the  large  amount  of  carbon  set 
free  explains  the  black  smoke  produced: 

2CH3C6H2(N03)3  ->  5H20  +  3N2  +  6C02  +  CO  +  7C. 

6.  Organic  compounds  of  another  class,  the  alcohols  (q.v.),  inter- 
act with  molecular  nitric  acid  in  a  different  way.  The  latter  is 
mixed  with  sulphuric  acid,  which  assists  in  the  removal  of  the  ele- 


350  COLLEGE    CHEMISTRY 

ments  of  water  (p.  286).  Thus,  when  glycerine  is  added  slowly  to 
the  cooled  mixture,  glyceryl  nitrate  (so-called  nitroglycerine)  is 
produced  : 

C3H5(OH)3  +  3HN03  -»  C3H5(N03)3  +  3H2O. 


Guncotton  is  made  by  this  action,  cotton  (cellulose)  being  em- 
ployed : 

(C6HioO6)2  +  6HNO3  ->  Ci2Hi4O4(NO3)6  +  6H2O. 


7.  Nitric  acid  produces  substances  of  bright-yellow  color,  known 
as  xanthoproteic  acids,  when  it  comes  in  contact  with  proteins,  e.g., 
in  the  skin,  or  in  wool.  Hence  nitric  acid  stains  woolen  clothing 
yellow.  This  reaction  is  used  as  a  test  for  proteins. 

Nitrates.  —  The  nitrates  are  all  more  or  less  easily  soluble  in 
water.  When  heated  they  decompose  in  one  or  other  of  three  ways 
(see  pp.  351,  356,  357).  The  individual  nitrates,  such  as  sodium 
nitrate  and  potassium  nitrate,  are  described  elsewhere. 

NITRIC  OXIDE  AND  NITROGEN  TETROXIDE 

Preparation  of  Nitric  Oxide  NO.  —  Pure  nitric  oxide  is  ob- 
tained by  adding  nitric  acid  to  a  boiling  solution  of  ferrous  sulphate 
in  dilute  sulphuric  acid  or  of  ferrous  chloride  in  hydrochloric  acid  : 

2FeS04  +  H2S04  ->  Fe2(S04)3  (+  2H)    X  3.         (1) 
(3H)  +  HN03  -»  NO  +  2H2O  X  2.        (2) 

6FeSO4  +  3H2SO4  +  2HNO3  -»  3Fes(SO4)8  +  2NO  +  4H2O. 

The  first  partial  equation  does  not  take  place  at  all  unless  an  oxi- 
dizing agent  like  nitric  acid  is  present  (p.  225).  The  multiplication 
of  the  two  partial  equations  by  3  and  2,  respectively,  is  required  in 
order  that  the  hydrogen,  which  is  not  a  product,  may  cancel  out. 
This  action  is  used  as  a  means  of  determining  the  quantity  of 
nitric  acid  in  a  solution,  or  of  nitrates  in  a  mixture,  by  measure- 
ment of  the  volume  of  nitric  oxide  evolved. 

As  we  shall  see  (p.  354),  nitric  oxide  may  also  be  obtained  when 
sufficiently  dilute  nitric  acid  (sp.  gr.  1.2)  acts  upon  copper.  This 
interaction  furnishes  the  most  convenient  method  of  generating 
the  gas  in  the  laboratory  (see  also  p.  352). 


OXIDES  AND   OXYGEN  ACIDS   OF  NITROGEN  351 

Properties  of  Nitric  Oxide.  —  Nitric  oxide  is  a  colorless  gas. 
In  solid  form  it  melts  at  —  167°,  and  the  liquid  boils  at  —  153.6°. 
Its  solubility  in  water  is  slight.  The  density  of  the  gas  shows  the 
formula  to  be  NO;  and  there  is  no  tendency  to  form  a  polymer, 
such  as  N2O2,  even  at  low  temperatures. 

This  gas  is  the  most  stable  of  the  oxides  of  nitrogen.  Vig- 
orously burning  phosphorus  continues  to  burn  in  the  gas,  nitrogen 
being  set  free.  Burning  sulphur  and  an  ignited  taper,  however, 
are  extinguished. 

Nitric  oxide  has  two  characteristic  chemical  properties.  It  unites 
directly  with  oxygen  in  the  cold  to  form  the  reddish-brown  nitrogen 
tetroxide  : 


The  same  result  follows  when  it  is  led  into  warm  concentrated 
nitric  acid:  NO  +  2HN03  <=^  3NO2  +  H20. 

It  also  unites  with  a  number  of  salts,  the  compound  in  the  ca^se  of 
ferrous  sulphate,  FeNO.SO4,  being  capable  of  existence  in  solution 
and  possessing  a  brown  color.  Since  ferrous  sulphate  will  first 
reduce  nitric  acid  to  nitric  oxide  (p.  350),  and  the  excess  of  the  salt 
will  then  give  a  brown  color  with  the  product,  a  delicate  test  for 
nitric  acid  is  founded  upon  these  actions. 

Preparation  of  Nitrogen  Tetroxide  NO2*  —  This  substance 
is  liberated  by  heating  nitrates,  other  than  those  of  potassium, 
sodium,  or  ammonium,  such  as  lead  and  copper  nitrates: 

2Cu(N03)2  ->  2CuO  +  4N02  +  02. 

The  oxide  of  the  metal  remains,  unless  this  oxide  is  itself  decom- 
posed by  heating  (p.  60).  When  the  mixed  gases  are  led  through 
a  U-tube  immersed  in  ice,  the  tetroxide  condenses  as  a  yellow 
liquid  (b.-p.  22°,  m.-p.  —10.5°),  and  the  oxygen  passes  on. 

The  compound  may  also  be  made  by  direct  union  of  nitric  oxide 
and  oxygen,  or  by  oxidation  of  nitric  oxide  by  concentrated  nitric 
acid  (p.  351).  It  is  likewise  almost  the  sole  product  of  the  inter- 
action of  concentrated  nitric  acid  with  tin  or  copper  (see  p.  355)  .  If 
any  nitric  oxide  were  produced  by  the  primary  action,  it  would 
be  oxidized  to  nitrogen  tetroxide  in  passing  up  through  the  acid 
(p.  351). 


352  COLLEGE    CHEMISTRY 

Properties  of  Nitrogen  Tetroxide.  —  The  most  striking 
peculiarity  of  this  gas  is  that,  when*  hot,  it  is  deep  brown  in  color, 
and  when  cold,  pale  yellow.  The  density  of  the  brown  gas,  at  140°, 
corresponds  to  the  formula  NO2,  that  of  the  yellow  gas  at  22°,  to 
N204.  When  the  temperature  is  carried  above  154°,  by  passing  the 
brown  gas  through  a  red-hot  tube,  the  brown  color  disappears,  and 
nitric  oxide  and  oxygen  are  formed.  On  cooling,  the  same  steps 
through  brown  gas  to  pale-yellow  gas  are  retraced: 

2NO  +  02  <=>  2N02  4=*  N204 

Colorless  Brown         Colorless 

Since  nitrogen  tetroxide  yields  free  oxygen  more  readily  than 
does  nitric  oxide,  phosphorus  burns  readily  in  it;  a  taper,  however, 
is  extinguished.  On  account  of  its  oxidizing  power,  it  is  sometimes 
used  in  bleaching  flour. 

This  oxide  is  intermediate  in  composition  between  nitrous  and 
nitric  anhydrides,  and,  when  dissolved  in  cold  water,  gives  both 
nitric  and  nitrous  acids:  N204  +  H20  ->  HNO3  +  HN02.  If  a 
base  is  present,  a  mixture  of  the  nitrate  and  nitrite  of  the  metal  is 
produced.  When  the  water  is  not  cooled,  the  nitrous  acid  (q.v.\ 
being  unstable,  gives  nitric  oxide  and  nitric  acid:  3N02  +  H2O 
?±  2HN03  +  NO. 

NITRIC  ACID  FROM  ATMOSPHERIC  NITROGEN.. 

The  Reactions  Involved.  —  Nitrogen  and  oxygen  have  no 
tendency  to  unite  at  room  temperature  to  form  nitric  oxide.  The 
union  is  endothermal,  and  is  therefore  favored  by  a  high  temper- 
ature (Van't  Hoff's  law,  p.  188) : 

N2  +  02  +  43,200  cal.  +±  2NO. 

Even  at  2000°,  however,  using  atmospheric  air,  only  1  per  cent  of 
nitric  oxide  is  formed,  and  at  3000°,  5  per  cent.  The  electric  dis- 
charge actually  used  gives  about  1  per  cent. 

The  mixture  is  next  cooled,  to  permit  the  union  of  2NO  +  02 
^  2N02,  because  (p.  352)  nitrogen  tetroxide  is  decomposed  at 
about  154°,  and  therefore  cannot  be  formed  at  2000°. 

Next,  the  air  containing  NO2  is  passed  through  absorbing  towers, 
down  which  water  trickles,  and  nitric  acid  is  formed: 

3N02  +  H20  ->  2HN03  +  NO. 


OXIDES   AND  OXYGEN  ACIDS    OF  NITROGEN 


353 


The  NO  liberated  combines  with  more  atmospheric  oxygen  to 
form  N02,  which  interacts  again  with  the  water,  and  practically 
no  nitric  oxide  is  lost. 

Finally,  the  nitric  acid  is  poured  upon  limestone  (CaC03),  and 
the  calcium  nitrate  formed  is  sold  for  use  as  a  fertilizer,  under  the 
name  air  saltpeter. 

The  Plant   used   in   the  Fixa- 
tion.—  At  Notodden  and  elsewhere 
WATER  f0a  \  in    Norway,    the 

Birkeland-Eyde 
process  (Fig.  96) 
is  used.  Hydro- 
electric power  is 
employed,  and  an 
arc  discharge  be- 
tween two  rods  of 
carbon  is  spread, 
by  the  influence  of 
large  and  powerful 

electromagnets,  into  a  circular  brush  discharge 
several  feet  in  diameter.  The  figure  is  a  cross 
section  of  the  space  filled  by  the  discharge,  the 
small  circle  in  the  center  being  a  section  of  one 
carbon  rod.  Air  is  blown  through  the  flame  in 
such  a  way  that  none  can  avoid  passing  through 
at  least  a  part  of  the  heated  area.  The  yield  is 
about  70  g.  of  nitric  acid  per  kilowatt-hour,  and 
the  net  earnings  are  $350,000  (1911). 

The    Badische    process,    used   in   the   same 
factories    in    Norway,   employs    a    discharge 
through  a  tube  over  20  feet  long  (Fig.  97). 
The  stream  of  air  rotates  as  it  traverses  the 
tube,  so  that  every  part  is  exposed  to  the  dis- 
charge.    The  Pauling  process,  used  at  Gelsen- 
kirchen  in  Germany  and  Nitrolee,  South  Caro- 
lina, uses  preheated  air  and  a  different  arrangement  of  the  discharge. 
For  other  reactions  involving  the  fixation  of  atmospheric  nitro- 
gen, see  calcium  cyanamide  (q.v.)  and  root  nodules  (p.  339). 


FIG.  96. 


\ENTRMCe 


FIG.  97. 


354  COLLEGE    CHEMISTRY 

OXIDIZING  ACTIONS  OF  NITRIC  ACID 

When  nitric  acid  gives  up  oxygen  to  any  body,  it  is  itself  reduced. 
Hence,  according  to  convenience,  we  shall  refer  to  oxidations  by, 
or  reductions  of  nitric  acid. 

Oxidation  of  Hydrogen.  —  The  metals  preceding  hydrogen 
in  the  electromotive  series  (p.  260)  displace  hydrogen  from  nitric 
acid,  as  they  do  from  other  acids.  With  metals  more  active  than 
zinc,  such  as  magnesium,  a  great  part  of  the  hydrogen  escapes  in 
the  free  condition.  But,  in  the  case  of  zinc  and  the  metals  below  it, 
most  or  all  of  the  hydrogen  is  oxidized  to  water  by  the  nitric  acid, 
and  part  of  the  acid  is  reduced  (see  Active  hydrogen,  p.  360). 
Thus,  with  zinc  and  very  dilute  nitric  acid,  almost  the  only  product, 
aside  from  zinc  nitrate,  is  ammonia: 

4Zn  +  8HN03  -*  4Zn(N03)2(+  8H).  (1) 

(8H)  +  HN03  ->  NH3  +  3H20.  (2) 

(3) 


4Zn  +  10HNO3^4Zn(N03)2  +  NILJSTO;,  +  3H2O. 
With  the  excess  of  nitric  acid  (3),  ammonium  nitrate  is  formed. 

Heavy  Metals.  —  The  less  active  metals,  such  as  copper  and 
silver,  do  not  displace  hydrogen  from  dilute  acids  (p.  60),  but 
reduce  nitric  acid,  nevertheless,  and  are  converted  into  nitrates. 
Platinum  and  gold  (cf.  p.  287)  alone  are  not  attacked.  Thus, 
copper,  with  somewhat  diluted  nitric  acid  (sp.  gr.  1.2),  gives  cupric 
nitrate  and  nitric  oxide  NO. 

In  making  the  equation  for  this  action  we  may  use  the  anhydride 
plan  (p.  325),  which  is  applicable  whenever  an  oxygen  acid  gives 
an  oxide  by  reduction.  We  resolve  the  formula  of  nitric  acid  into 
those  of  water  and  the  anhydride  H20,N2O5(  =  2HNO3).  This 
shows  that  the  two  molecules  of  the  acid  will  give  2NO,  and  3O 
will  remain: 

2HN03  (or  H2O,N205)  ->  H20  +  2NO  (+  30).  (1) 

(3O)  +  6HNO3  +  3Cu  -*  3H20  +  3Cu(NO3)2.  (2) 

8HN03  +  3Cu  ->4H2O  +  2NO  +  3Cu(NO3)2. 

The  nitric  oxide  is  liberated  as  a  colorless  gas,  but  forms  the  brown 
tetroxide  at  once  on  meeting  the  oxygen  of  the  air  (p.  351). 


OXIDES  AND   OXYGEN  ACIDS   OF  NITROGEN  355 

When  concentrated  nitric  acid  is  used  with  copper,  almost  pure 
nitrogen  tetroxide  is  obtained: 

2HNO3  (or  H2O,N205)  -»  H20  +  2N02  (+  0).  (1) 

(O)  +  2HNOS  +  Cu  ->  H20  +  Cu(N03)2. (2) 

4HNO3  +  Cu  -»  2H20  +  2NO2  +  Cu(N03)2. 

The  reader  should  note  the  constant  production  of  nitric  oxide  with 
diluted  nitric  acid,  and  the  invariable  formation  of  nitrogen  tetroxide 
with  concentrated  acid.  This  is  explained  by  the  fact  that  nitrogen 
tetroxide  cannot  pass  unchanged  through  a  liquid  containing  much 
water,  for  it  gives  nitric  acid  and  nitric  oxide  with  the  latter  (p.  352). 
Conversely,  where  the  nitric  acid  is  concentrated,  nitric  oxide,  even 
if  formed  by  the  interaction  with  the  metal,  must  be  oxidized  to 
nitrogen  tetroxide  as  it  passes  up  through  the  liquid  (p.  351). 
Note,  also,  that  the  nitrate  of  the  metal  is  formed,  if  the  nitrate  is 
not  hydrolyzed  by  water,  not  the  oxide. 

Oxidation  of  Non-Metals.  —  With  non-metals  the  actions  are 
different,  in  so  far  that  these  elements  form  no  nitrates.  Thus 
sulphur  boiled  in  nitric  acid  gives  sulphuric  acid,  along  with  nitric 
oxide,  equation  (3),  or  with  nitrogen  tetroxide,  equation  (6),  or 
with  both,  according  to  the  concentration  of  the  acid  (see  above)  : 

2HN03  (or  H2O,N205)  ->  2NO  +  H20  (+  30).  (1) 

(3O)  +  H2O  +  S  -*  H2SO4. (2) 

2HNO3  +  S  -+  2NO  +  H2S04.  (3) 

2HN03  (or  H2O,N205)  -»  2N02  +  H20  +  0       X  3.  (4) 

(3O)  +  H20  +  S  -*  H2SO4. <  (5) 

6HN03  +  S  -»  6N02  +  2H20  +  H2S04.  (6) 

The  reader  will  note  that  a  separate  equation,  (3)  and  (6),  must 
be  made  for  the  formation  of  each  reduction  product.  If  NO  and 
N02  are  both  formed,  they  cannot  arise  from  the  same  molecule 
of  nitric  acid.  They  result  from  two  actions  which  are  independent, 
although  proceeding  concurrently  in  the  same  vessel  (cf.  p.  317). 
Thus  the  equation:  2HN03  +  C  -»  H2O  +  C02  +  NO  +  N02,  is 
a  misrepresentation.  It  implies  that  equimolar  quantities  of  the 
two  oxides  of  nitrogen  are  formed.  But  this  could  occur  only  by 
chance,  and  the  balance  would  be  destroyed  the  next  moment  by 


356  COLLEGE    CHEMISTRY 

the  lowering  in  the  concentration  of  the  acid,  giving  the  advantage 
to  the  nitric  oxide. 

Oxidation  of  Compounds:  Aqua  Regia.  —  Compounds  like 
hydrogen  sulphide  and  sulphurous  acid,  which  are  easily  oxidized, 
interact  with  nitric  acid.  With  diluted  nitric  acid,  the  products  are 
free  sulphur  and  sulphuric  acid  respectively. 

The  mixture  of  nitric  acid  and  hydrochloric  acid  is  known  as 
aqua  regia.  Chlorine  is  set  free  by  the  oxidation  of  the  hydro- 
chloric acid, 

HN03  +  3HC1  -»  2H20  +  Ct  +  NOC1, 

and  nitrosyl  chloride  NOC1  is  also  formed.  The  liquid  thus  con- 
tains several  oxidizing  agents,  nitric  acid,  hypochlorous  acid  (from 
C12  +  H20),  and  some  nitrous  acid.  It  is  frequently  used  in 
analysis,  for  example  to  oxidize  sulphur  (say,  in  cast  iron  or  in 
minerals),  the  sulphuric  acid  formed  being  estimated  by  precipi- 
tation and  weighing  of  barium  sulphate  (p.  287). 

Aqua  regia  (Lat.,  royal  water)  received  its  name  because  it  con- 
verted the  "noble"  metals,  gold  and  platinum,  into  soluble  com- 
pounds. This  it  does  because  the  free  chlorine,  in  presence  of 
hydrochloric  acid,  combines  to  form  the  exceedingly  stable  com- 
plex ions  (see  pp.  505, 508)  AuCLj"  (see  chlorauric  acid,  and  PtCl6=, 
the  negative  ion  of  chloroplatinic  acid: 

2HC1  +  2C12  +  Pt  -»  H2PtCl6,  or  Pt  +  2C12  +  2C1~  -» RC18=. 

NITROUS  ACID,  HYPONITROUS  ACID,  AND  THEIR  ANHYDRIDES 

Nitrites  and  Nitrous  Acid.  —  When  the  nitrates  of  potassium 
and  sodium  are  heated,  they  lose  one  unit  of  oxygen,  and  the 
nitrites  remain: 

2NaN03  -» 2NaN02  +  02. 

Commonly  lead  is  stirred  with  the  melted  nitrate  and  assists  in  the 
removal  of  the  oxygen.  The  litharge  PbO  which  is  formed  re- 
mains as  a  residue  when  the  sodium  nitrite  is  dissolved  for  re- 
crystallization. 

When  an  acid  is  added  to  a  dilute  solution  of  a  nitrite,  a  pale-blue 
solution  containing  nitrous  acid  HN02  is  obtained.  The  acid  is 


OXIDES   AND  OXYGEN  ACIDS   OF  NITROGEN  357 

very  unstable,  however,  and,  when  the  solution  is  warmed,  it  de- 
composes : 

3HN02  -» HN03  +  2NO  +  H2O. 

When  a  concentrated  solution  of  sodium  nitrite  (or  the  solid  salt 
itself)  is  acidified,  the  nitrous  acid  decomposes  at  once,  and  a  brown 
gas  containing  the  anhydride  escapes: 

2H+  +  2N02~  fc?  2HNO2  fc?  H20  +  N203T . 

This  behavior  distinguishes  a  nitrite  from  a  nitrate. 
Nitrous  acid  is  an  active  oxidizing  agent: 

2HI  +  2HN02  (or  H2O,N203)  -» 2H20  +  2NO  +  I2. 

Indigo  is  also  converted  by  it  into  isatin  (cf.  p.  311). 
Nitrous  acid  is  much  used  in  the  making  of  organic  dyes. 

Nitrous  Anhydride  N2O3.  —  A  study  of  the  density  of  the  gas 
arising  from  the  decomposition  of  nitrous  acid  shows  that,  in  the 
gaseous  state,  the  anhydride  is  almost  entirely  dissociated: 

N203  ^±  NO  +  N02. 

When  the  mixture  is  led  through  a  U-tube  immersed  in  a  freezing 
mixture  at  —21°,  a  deep-blue  liquid  is  obtained  which  is  the 
anhydride  itself.  This  dissociates  rapidly  when  allowed  to  boil. 

The  same  equimolar  mixture  of  the  two  gases  is  obtained  by  the 
action  of  water  on  nit  rosy  Isulphuric  acid  (p.  281). 

Hyponitrous  Acid  and  Nitrous  Oxide  N2O.  —  Hyponitrous 
acid  H2N202  is  a  white  solid.  Its  solution  in  water  is  an  exceed- 
ingly feeble  acid.  The  warm  aqueous  solution  decomposes  slowly, 
giving  nitrous  oxide: 

H2N2O2  ->  H2O  +  N2O, 

and  this  change  is  not  capable  of  reversal. 

Nitrous  oxide  is  prepared  by  gently  heating  ammonium  nitrate 
(an  explosive),  or  a  solution  of  a  salt  of  ammonium  and  a  nitrate: 

NH4+  +  NOS~  ^  NH4NO3  -4  2H2O  +  N20. 

The  steam  condenses,  and  the  nitrous  oxide  may  be  collected  over 
warm  water,  or  be  dried  and  compressed  into  steel  cylinders. 


358  COLLEGE    CHEMISTRY 

Its  solubility  in  cold  water  is  considerable  :  '  at  0°,  130  volumes  in 
100;  at  25°,  60  in  100.  The  liquefied  gas  boils  at  -89.8°  and  its 
vapor  tension  at  20°  is  49.4  atmospheres. 

A  glowing  splinter  of  wood  bursts  into  flame  in  nitrous  oxide, 
and  phosphorus  and  sulphur  burn  in  it  with  much  the  same  vigor 
as  in  oxygen.  In  all  cases  oxides  are  formed,  and  nitrogen  is  set 
free.  It  does  not  combine  with  nitric  oxide,  however,  as  does 
oxygen  (p.  351). 

Metals  do  not  rust  in  nitrous  oxide,  and  the  haemoglobin  of  the 
blood  is  unable  to  use  it  as  a  source  of  oxygen.  It  is  employed  as  an 
anaesthetic  for  minor  operations.  The  hysterical  symptoms  which 
accompany  its  use  caused  it  to  receive  the  name  of  "laughing  gas." 

Graphic  Formulae  of  Nitric  Acid  and  its  Derivatives:  Ex- 
plosives. —  The  following  equation  for  the  formation  of  am- 
monium nitrate  by  neutralization  of  ammonium  hydroxide  with 
nitric  acid  shows  the  graphic  (p.  292)  or  structural  formulae  of 
these  substances: 


H/  H/ 

The  structural  formula  of  the  nitrate  is  intended  to  explain  the  fact 
that  the  salt  is  able  to  exist  at  all,  by  representing  the  oxygen  and 
hydrogen  as  being  separated  from  one  another  and  attached  to 
different  nitrogen  units.  When  the  equilibrium  of  the  system  is 
disturbed  by  heating,  the  oxygen  and  hydrogen  unite  to  form 
water,  an  arrangement  which  is  much  more  stable,  and  nitrous 
oxide  (p.  357)  escapes  with  the  steam. 

The  behavior  of  nitroglycerine  and  guncotton  (p.  350),  as  well 
as  of  ammonium  nitrite  (p.  338),  is  explained  in  the  same  way. 
These  substances  are  made  by  actions  which,  like  the  above 
neutralization,  take  place  in  the  cold,  and  the  groups,  containing  the 
oxygen  on  the  one  hand  and  carbon  and  hydrogen  on  the  other, 
become  quietly  united  without  more  serious  interaction.  When, 
however,  the  nitroglycerine,  for  example,  is  heated,  or  receives  a 
mechanical  shock,  the  carbon  and  hydrogen  unite  with  the  oxygen 
and  the  nitrogen  escapes: 

4C3H5(N03)3  ->  12C02  +  10H20  +  6N2  +  02. 


OXIDES  AND   OXYGEN  ACIDS   OF  NITROGEN  359 

Smokeless  Powder  and  Dynamite.  —  Dried  guncotton  (p. 
350)  simply  burns  briskly  (deflagrates)  when  set  on  fire.  Whether 
wet  or  dry,  it  explodes,  but  only  from  a  suitable  shock,  such  as 
that  produced  by  fulminate  of  mercury  Hg(OCN)2,  used  in  per- 
cussion caps.  In  pure  form  it  is  used  only  in  torpedoes  or  sub- 
marine mines.  Like  nitroglycerine  (p.  350),  it  explodes  too 
rapidly,  and  would  burst  the  gun,  or  pulverize  the  ore  or  coal  if 
used  for  blasting.  Neither  of  these  substances  "explodes  down- 
wards only."  The  explosion  strikes  the  air  with  equal  violence, 
but  the  effect  on  the  air  escapes  notice  because  it  is  not  permanent, 
while  the  shattering  of  a  rock  or  plate  of  steel  remains. 

Cordite,  one  variety  of  smokeless  powder,  is  made  by  dissolving 
guncotton  (65  parts),  nitroglycerine  (30  parts)  and  vaseline  (5 
parts)  in  acetone.  The  resulting  paste  is  rolled  out  and  cut  into 
small  pieces.  When  the  acetone  evaporates,  the  horny  cordite 
remains.  These  explosives  are  smokeless  because,  unlike  gun- 
powder and  T.N.T.,  they  yield  no  solids  when  they  decompose 
(see  equations). 

Various  forms  of  dynamite  are  made  like  cordite,  excepting  that 
sodium  or  ammonium  nitrate  and  sawdust  or  flour  are  added,  so 
that  the  rate  of  explosion  may  be  regulated  and  the  coal  or  ore  may 
be  split  up,  but  not  shattered  or  pulverized. 

Plastics.  —  A  guncotton,  less  completely  " nitrated"  by  nitric 
acid,  when  worked  between  rollers  with  camphor  and  a  little 
alcohol,  gives  a  viscous  solution  (Parkes,  before  1855).  When  the 
alcohol  evaporates,  transparent,  colorless  celluloid  (first  made  by 
Hyatt)  remains.  The  moist  dough  is  rolled  into  sheets  to  make 
photographic  films.  By  adding  dyes  and  "fillers,"  and  molding 
the  dough,  black  combs,  brush  handles,  white  knife  handles,  etc., 
can  be  manufactured. 

Collodion  is  a  solution  of  the  same  guncotton  in  a  mixture  of 
alcohol  and  ether.  When  collodion  is  forced  through  minute  holes 
in  a  steel  dye,  the  threads  dry  as  they  come  out  and  can  be  wound 
on  spools.  Treatment  with  an  alkali  "  denitrates "  the  threads, 
restoring  the  composition  to  that  of  the  original  cotton.  The  prod- 
uct, one  of  the  forms  of  artificial  silk,  is  at  least  as  brilliant  as  the 
real  article  (a  protein,  not  related  chemically  to  cotton),  and  sus- 
ceptible of  being  dyed  to  any  desired  tint. 


360  COLLEGE   CHEMISTRY 

Balancing  Equations.  —  The  reader  should  practice  the 
balancing  of  the  equations  for  oxidations  occurring  in  this  chapter, 
using  all  the  methods.  We  have  used  the  anhydride  plan  (p.  355) 
and  that  of  partial  equations  (p.  350).  To  illustrate  the  other 
plans,  take  the  action  of  concentrated  nitric  acid  on  copper  (p. 
355). 

Positive  and  negative  valence  method  (p.  322)  .  Write  the  skeleton 
equation  : 

Skeleton:        HN03  +  Cu  -»  H20  +  N02  +  Cu(N03)2. 

We  perceive  that  on  the  left  the  valence  of  03  is  —6  and  of  H  is 
+  1  :  that  of  N  is  therefore  +5.  That  of  Cu  is  zero.  On  the  right, 
the  valence  of  N  is  +4  and  of  Cu  +2.  Evidently,  2N  changing 
from  2  X  +5  to  2  X+4  will  furnish  +2  for  the  copper.  We 
note  also  that  2NO3  is  required,  without  change,  for  Cu(NO3)2. 
Hence,  altogether  4HNO3  is  needed  on  the  left,  and  gives  2NO2: 

Balanced:    4HN03  +  Cu  -»  2H2O  +  2N02  +  Cu(NO3)2. 

Positive  electrical  charge  plan  (p.  325).  In  the  skeleton  equation 
(above)  we  first  separate  the  oxidizing  ions  and  their  products  from 
the  oxidized  substance  and  the  change  it  undergoes: 

NOr  +  2H+         ->  N02°  +  H2O  +  0       X  2. 
Cu°  +  20 


_  _ 

Cu°  +  2NO3~  +  4H+  ->  2NO2°  +  2H2O  +  Cu++. 

The  first  partial  equation  produces  ©,  while  the  second  requires 
2  ©,  and  hence  the  former  is  multiplied  by  2  before  the  addition 
takes  place.  Since  NOs~  is  the  only  acid  radical  present,  it  is 
understood  that  cupric  nitrate  is  the  salt  formed. 

Active  ("Nascent")  Hydrogen.  —  When  hydrogen  gas  is  led 
through  cold  nitric  acid,  little  or  no  action  occurs.  But  (p.  354) 
when  zinc,  which  interacts  with  acids  to  give  hydrogen,  is  placed 
in  nitric  acid  the  latter  is  reduced.  To  explain  the  apparent 
greater  activity  of  the  hydrogen  in  the  second  instance,  we  note 
the  fact  that  it  is  liberated  on  the  surface  of  the  zinc.  The  contact 
(catalytic)  effect  of  the  zinc  increases  its  activity.  Many  metals 
have,  in  a  greater  or  less  degree,  this  power  of  increasing  the 
activity  of  hydrogen.  Thus,  hydrogen  absorbed  in  platinum  or 


OXIDES  AND    OXYGEN  ACIDS   OF  NITROGEN  361 

palladium  (p.  57)  or  liberated  by  electrolysis  on  poles  made  of 
these  metals,  reduces  nitric  acid  readily.  Other  elements,  such  as 
the  oxygen  used  in  making  sulphur  trioxide  (p.  279),  are  also  ren- 
dered more  active  by  contact  agents. 

This  more  active  state  of  hydrogen  is  described  as  the  nascent 
state,  because  it  happens  to  be  a  common  condition  of  hydrogen 
when  associated  with  substances  which  produce  it.  The  active 
state  has,  however,  no  necessary  connection  with  such  an  immedi- 
ately preceding  act  of  liberation,  as  the  platinum  and  sulphur 
trioxide  illustrations,  and  the  following  experiment  show:  Three 
test-tubes  are  filled  with  dilute,  acidified  potassium  permanganate 
solution.  Zinc  dust,  added  to  one,  generates  hydrogen  and  causes 
decolorization.  A  little  platinum  black  is  added  to  the.  second,  and 
hydrogen  gas  is  led  through  this  and  the  third.  The  contact 
action  of  the  platinum  enables  the  hydrogen  quickly  to  reduce 
the  permanganate,  while  the  third  portion  remains  unaltered. 

Besides,  if  the  action  were  due  to  freshly  liberated,  perhaps 
atomic  hydrogen,  this  substance  should  have  constant  properties. 
But  it  has  not.  Thus,  nitric  acid  with  zinc  gives  much  ammonia; 
with  magnesium,  none;  with  tin,  ammonia  and  hydroxylamine 
HONH2  as  well. 

Exercises.  —  1.  Make  the  equation  for  the  interaction  of 
ferrous  chloride,  hydrochloric  acid,  and  nitric  acid  (p.  350),  and 
for  all  the  actions  concerned  when  the  test  for  a  nitrate  (p.  351)  is 
applied  to  sodium  nitrate.  What  volume  (at  0°  and  760  mm.)  of 
NO  is  obtained  from  one  formula-weight  of  nitric  acid  (p.  350)? 

2.  In  the  action  of  zinc  on  dilute  nitric  acid  (p.  354),  why  is  not 
the  ammonia  given  off  as  a  gas?     How  should  you  show  that  it  was 
formed  at  all? 

3.  Make  correct  equations  for  the  formation  of  nitric  oxide  and 
nitrogen  tetroxide  by  the  action  of  carbon  on  nitric  acid  (p.  355). 

4.  Make  equations  for  the  interaction  of  iron  with  diluted  and 
with  concentrated  nitric  acid,  respectively  (p.  355) .     The  iron  gives 
ferric  nitrate  Fe(NO3)3. 

5.  Give  the  three  ways  in  which  nitrates  decompose  when 
heated,  with  one  equation  illustrating  each. 

6.  Make  all  the  equations  for  oxidations  on  pp.  350  and  354, 
using  the  methods  illustrated  on  p.  360. 


CHAPTER  XXVII 
PHOSPHORUS 

The  Chemical  Relations  of  the  Element.  —  There  are 
many  things  in  the  chemistry  of  phosphorus  and  its  compounds 
which  remind  us  of  nitrogen.  Yet  these  are  largely  referable  to 
the  fact  that  the  elements  are  both  non-metals  and  both  have  the 
same  valences,  viz.,  three  and  five.  The  behavior  of  the  com- 
pounds is  often  very  different.  For  the  present  it  is  sufficient  to 
say  that  both  give  compounds  with  hydrogen,  NH3  and  PH3,  and 
both  yield  oxides  of  the  forms  X2O3,  X2O4,  and  X2O5.  The  first 
and  last  of  these  oxides  are  acid-forming,  and  phosphorus,  there- 
fore, gives  acids  corresponding  to  nitrous  and  nitric  acids.  The 
element  is  thus  a  non-metal. 

Occurrence.  —  This  element  is  found  in  nature  in  the  form  of 
phosphates.  Calcium  phosphate  CaaCPO^  forms  26  per  cent  of 
the  bones  and  teeth,  and  it  occurs  in  all  fertile  soils.  It  consti- 
tutes a  large  part  of  the  "phosphate  rock"  of  Georgia,  Florida, 
the  Carolinas,  Tennessee,  and  of  Algeria  and  Tunis.  A  con- 
spicuous mineral  related  to  this  substance,  apatite,  Ca5F(P04)3 
and  Ca6Cl(PO4)3,  is  found  in  large  quantities  in  Canada,  and  is 
a  component  of  many  rocks.  Complex  organic  compounds  of 
phosphorus,  such  as  lecithin,  are  essential  constituents  of  the 
muscles,  the  nerves,  and  the  brain.  Amongst  foods,  egg-yolks  and 
beans  contain  a  large  proportion. 

Preparation.  —  Brand,  merchant  and  alchemist,  of  Hamburg, 
discovered  phosphorus  (1669)  by  distilling  the  residue  from  evapo- 
rated urine,  in  the  course  of  his  search  for  the  philosopher's  stone. 
The  mode  of  preparing  it  from  bone-ash,  which  contains  83  per 
cent  of  calcium  phosphate,  was  first  published  by  Scheele  (1771). 
Now  the  less  expensive  calcium  phosphate  of  fossil  origin  is 
employed. 

362 


PHOSPHORUS 


363 


The  calcium  phosphate  is  mixed  with  the  proper  proportions  of 
carbon  and  silicon  dioxide  (sand),  and  the  mixture  is  introduced 
continuously  into  an  electric  furnace  (Fig.  98).  The  discharge  of 
an  alternating  current  between  carbon  poles 
produces  the  very  high  temperature  which 
the  action  requires.  The  calcium  silicate 
which  is  formed  fuses  to  a  slag,  and  can  be 
withdrawn  at  intervals.  The  gaseous  prod- 
ucts pass  off  through  a  pipe  and  the  phos- 
phorus is  caught  under  water: 

Ca3(P04)2  +  3Si02  +  5C  ->  3CaSi03 
+  5CO  +  2P. 


FIG.  98. 


We  may  regard  the  phosphate  as  being  com- 
posed of  two  oxides,  3CaO,P2O5.      It  thus 
appears  that  the  calcium  oxide  has  united 
with  the  silica,  which  is  an  acid  anhydride  (cf.  p.  280) :  CaO  +  SiO2 
— >  CaSi03,  while  the  phosphoric  anhydride  has  been  reduced. 

The  phosphorus,  after  purification,  is  cast  into  sticks  in  tubes  of 
tin  or  glass,  standing  in  cold  water. 

The  Electric  Furnace.  —  By  an  electric  furnace  is  understood 
an  electrothermal  arrangement  in  which  the  heat  produced  by  some 
resistance  offered  to  the  current,  such  as  that  of  an  air-gap 
between  the  carbons,  is  used  to  produce  chemical  changes  re- 
quiring a  high  temperature.  Electrolysis  plays  no  part  in  the 
phenomena,  and  an  alternating  current,  which  can  produce  no 
electrolytic  decomposition,  is  generally  employed.  The  restricted 
area  within  which  the  heat  is  developed  makes  possible  the  attain- 
ment of  a  high  temperature  (see  Calcium  carbide). 

Physical  Properties.  —  There  are  at  least  two  allotropic  forms 
(p.  222)  of  phosphorus,  known  as  white  phosphorus  and  red 
phosphorus.  White  phosphorus,  prepared  as  described  above,  is  at 
first  transparent  and  colorless,  but  after  exposure  to  light  acquires 
a  superficial  coating  of  the  red  variety.  It  melts  at  44°  and  boils 
at  287°.  Its  sp.  gr.  is  1.83.  Its  molecular  weight  at  313°  is  128 
and  the  formula,  therefore,  P4  (cf.  p.  111).  Yellow  phosphorus 
is  soluble  in  carbon  bisulphide,  less  soluble  in  ether,  and  insoluble 


364  COLLEGE    CHEMISTRY 

in  water.  It  is  exceedingly  poisonous,  less  than  0.15  g.  being  a 
fatal  dose,  and  is  an  ingredient  in  roach-paste  and  rat  poison. 
Continued  exposure  to  its  vapor  causes  necrosis,  a  disease  from 
which  match-makers  are  liable  to  suffer.  The  jawbones  and  teeth 
are  particularly  liable  to  attack. 

Red  phosphorus  is  a  red  powder  consisting  of  small  tabular 
crystals.  It  is  obtained  by  heating  yellow  phosphorus  to  about 
250°  in  a  vessel  from  which  air  is  excluded.  Much  heat  is  evolved 
in  the  transformation.  Red  phosphorus  does  not  melt,  but 
passes  directly  into  vapor,  identical  with  that  of  yellow  phos- 
phorus. It  is  insoluble  in  carbon  bisulphide  and  other  solvents. 
It  is  not  poisonous,  and,  unlike  yellow  phosphorus,  does  not  re- 
quire to  be  kept  under  water  to  avoid  spontaneous  combustion. 
Red  phosphorus  appears  to  be  a  solid  solution  (p.  122)  of  the 
white  variety  in  a  less  active  kind.  Hence,  its  properties  are 
variable,  e.g.,  sp.  gr.  from  2.05  to  2.34.  Bridgeman,  by  heating 
white  phosphorus  at  200°  under  a  pressure  of  1200  kg./cm2.,  has 
obtained  black  phosphorus  (sp.  gr.  2.69)  which  may  be  the  pure 
form  of  the  red  variety. 

Chemical  Properties.  —  White  phosphorus  unites  directly 
with  the  halogens  with  great  vigor.  It  unites  slowly  with  oxygen 
in  the  cold,  and  with  sulphur  and  many  metals  when  the  materials 
are  heated  together.  The  slow  union  of  cold  phosphorus  with 
atmospheric  oxygen  is  accompanied  by  the  evolution  of  light. 
Hence  the  word  phosphorescence.  The  name  of  the  element  (Gk. 
</><os,  light;  <£e/>w,  I  bear)  records  this  property.  Apparently  the 
chemical  energy,  transformed  in  connection  with  the  oxidation, 
is  converted,  in  part  at  least,  into  radiant  energy  instead  of  com- 
pletely into  heat.*  The  slow  oxidation  of  phosphorus  is  ac- 
companied by  the  production  of  ozone,  but  the  nature  of  the 
action  is  still  unknown  (cf.  p.  219). 

Red  phosphorus,  since  it  is  formed  with  evolution  of  heat,  con- 
tains less  energy  than  white  phosphorus  and  is  much  less  active. 
It  does  not  catch  fire  below  240°,  while  ordinary  phosphorus 
ignites  at  35°. 

*  The  same  production  of  light  from  chemical  action  in  a  cold  body  is 
seen  in  the  luminosity  of  certain  parts  of  fireflies  and  some  species  of 
fish. 


PHOSPHORUS  365 

Matches.  —  In  making  common  matches,  invented  in  1827, 
the  splints  are  first  dipped  in  melted  sulphur  or  paraffin  to  the 
extent  of  about  half  an  inch.  The  head  is  often  composed  of 
lead  dioxide  Pb02,  which  supplies  oxygen,  a  small  proportion  of 
free  phosphorus  or  a  sulphide  of  phosphorus  P4S3  which  is  readily 
ignited  by  friction,  and  dextrin  or  glue.  The  use  of  white  phos- 
phorus is  forbidden  by  law  in  Sweden,  France,  Great  Britain  and 
Switzerland,  and  is  prevented  by  a  tax  of  two  cents  per  100 
matches  in  the  United  States. 

In  the  case  of  " safety"  matches,  the  mixture  upon  the  head  is 
not  easily  ignited  by  itself.  It  is  composed  of  potassium  chlorate 
or  dichromate,  some  sulphur  or  antimony  trisulphide  Sb2S3  (com- 
bustible), and  a  little  powdered  glass  or  chalk  to  increase  the 
friction,  all  held  together  with  glue.  Upon  the  rubbing  surface 
on  the  box  is  a  thin  layer  of  antimony  trisulphide  mixed  with  red 
phosphorus,  chalk  or  glass,  and  glue.  The  friction  converts  a 
little  of  the  red  phosphorus  into  vapor,  which  catches  fire  readily. 

Phosphine.  —  Three  hydrides  of  phosphorus  are  known. 
These  are,  phosphine  PH3  (a  gas),  a  liquid  hydride  P2H4,  which  is 
presumably  the  analogue  of  hydrazine  (N2H4),  and  a  solid  hydride 
P4H2. 

Phosphine  PH3  is  formed  slowly  by  the  action  of  active  hydrogen, 
from  zinc  and  hydrochloric  acid  at  70°,  upon  white  phosphorus. 
The  gas  may  be  made  by  boiling  white  phosphorus  with  potassium 
hydroxide  solution  in  a  flask  provided  with  a  delivery  tube. 
Potassium  hypophosphite  is  formed  at  the  same  time: 

3KOH  +  4P  +  3H20  ->  3KH2P02  +  PH3 1 . 

The  gas  made  in  this  way  contains  a  little  of  the  vapor  of  the  liquid 
hydride,  which  is  spontaneously  inflammable,  and  consequently  the 
bubbles  of  the  mixture  catch  fire  when  they  reach  the  surface  of 
water  in  the  trough:  PH3  +  2O2  ->  H3P04.  In  still,  moist  air, 
the  fog  of  droplets  of  phosphoric  acid  solution  form  smoke 
rings. 

The  simplest  method  of  preparing  phosphine  is  by  the  action  of 
water  upon  calcium  phosphide: 

Ca3P2  +  6H2O  ->  3Ca(OH)2  +  2PH3. 


366  COLLEGE    CHEMISTRY 

This  action  is  analogous  to  that  of  water  upon  magnesium  nitride 
(p.  339),  by  which  ammonia  is  produced.  In  consequence  of  the 
fact  that  calcium  phosphide  is  a  substance  of  irregular  compo- 
sition, a  mixture  of  all  three  hydrides  is  generally  obtained.  By 
passing  the  gas  through  a  strongly  cooled  delivery  tube,  however, 
the  liquid  and  solid  compounds  are  condensed  and  fairly  pure 
phosphine  passes  on. 

Phosphine  is  a  colorless  gas,  which  is  easily  decomposed  by  heat 
into  its  elements.  It  is  exceedingly  poisonous  and,  unlike  am- 
monia, it  is  insoluble  in  water,  and  produces  no  basic  compound 
corresponding  to  ammonium  hydroxide  when  brought  in  contact 
with  this  substance.  It  resembles  ammonia,  formally  at  least,  in 
uniting  with  the  hydrogen  halides  (see  below).  It  differs  from 
ammonia,  however,  inasmuch  as  it  does  not  unite  with  the  oxygen 
acids.  Phosphine  acts  upon  solutions  of  some  salts,  precipitating 
phosphides  of  the  metals: 

3CuSO4  +  2PH3  ->  Cu3P2 1  +  3H2S04. 

Phosphonium  Compounds.  —  Hydrogen  iodide  unites  with 
phosphine  to  form  a  colorless  solid,  crystallizing  in  beautiful, 
highly  refracting,  square  prisms:  PH3  +  HI  — >  PHJ.  Hydrogen 
chloride  combines  similarly  with  phosphine,  but  only  when  the 
gases  are  cooled  by  a  freezing  mixture,  or  are  brought  together 
under  a  total  pressure  of  18  atmospheres  at  14°.  When  the 
pressure  is  released,  rapid  dissociation  occurs. 

In  imitation  of  the  ammonia  nomenclature,  these  substances 
are  called  phosphonium  iodide  and  phosphonium  chloride  PHtCl. 
They  are  entirely  different,  however,  from  the  corresponding  am- 
monium derivatives,  for  the  PH4+  ion  is  unstable.  When  brought 
in  contact  with  water  they  decompose  into  their  constituents,  the 
hydrogen  halide  going  into  solution,  and  the  phosphine  being 
liberated  as  a  gas. 

Halides  of  Phosphorus.  —  The  existence  of  the  following 
halides  has  been  proved  conclusively: 

....  P2I4  (solid) 

PF3  (gas)        PC13  (liquid)         PBr3  (liquid)        PI3  (solid) 
PF6(gas)        PC15  (solid)  PBr5  (solid)          


PHOSPHORUS  367 

These  substances  may  all  be  formed  by  direct  union  of  the  elements. 
They  are  incomparably  more  stable  than  are  the  similar  com- 
pounds of  nitrogen.  They  are  all  hydrolyzed  by  water,  and  give 
an  oxygen  acid  of  phosphorus  and  the  hydrogen  halide  (see  below). 
This  action  was  used  in  the  preparation  of  hydrogen  bromide 
(p.  197)  and  hydrogen  iodide  (p.  201). 

Phosphorus  trichloride  PC13  is  made  by  passing  chlorine  gas  over 
melted  phosphorus  in  a  flask  until  the  proper  gain  in  weight  has 
occurred.  The  substance,  which  is  a  liquid  boiling  at  76°,  is  stable 
(cf.  p.  93).  When  excess  of  chlorine  is  employed,  phosphorus  pen- 
tachloride  PC15,  which  is  a  white  solid  body,  is  formed.  When 
moist  air  is  blown  over  any  of  these  substances,  the  water  is  con- 
densed to  a  fog  by  the  hydrogen  halide.  In  the  case  of  the  inter- 
action of  phosphorus  pentachloride  and  water,  phosphoric  acid  is 
formed : 

PC15  +  4H20  ~»  H3P04  +  5HC1. 

Phosphorus  pentachloride,  when  heated,  reaches  a  vapor 
tension  of  760  mm.  at  163°,  and  while  still  solid.  At  higher 
pressure  it  melts  at  166°.  At  these  temperatures,  about  4  per  cent 
of  the  molecules  are  dissociated  into  phosphorus  trichloride  and 
chlorine  (p.  117):  PC15  +±  PC13  +  C12. 

Oxides  of  Phosphorus.  —  The  oxides  of  phosphorus  are  the 
so-called  trioxide  P^e,  the  pentoxide  P205,  and  a  tetroxide  P204. 

The  pentoxide  is  a  white  powder  formed  when  phosphorus  is 
burned  with  a  free  supply  of  oxygen.  It  unites  with  water  with 
great  violence  to  form  metaphosphoric  acid  (see  below),  and  hence 
is  known  as  phosphoric  anhydride:  P205  +  H2O  — •»  2HP03.  In  the 
laboratory  this  action  is  frequently  utilized  for  drying  gases 
(p.  330)  and  for  removing  water  from  combination  (p.  349).  The 
vapor  density  leads  to  the  formula  P4Oi0,  use  of  which,  however, 
would  only  complicate  our  equations. 

The  trioxide  P^e  is  obtained  by  burning  phosphorus  in  a  tube 
with  a  restricted  supply  of  air.  It  is  a  white  solid,  melting  at 
22.5°  and  boiling  at  173°.  This  oxide  is  the  anhydride  of  phos- 
phorous acid,  but  it  unites  exceedingly  slowly  with  cold  water  to 
form  this  substance.  It  interacts  vigorously  with  hot  water,  but 
phosphine,  red  phosphorus,  hypophosphoric  acid,  and  phosphoric 
acid  are  amongst  the  products,  and  very  little  phosphorous  acid 


368  COLLEGE    CHEMISTRY 

escapes  decomposition.     When  this  oxide  is  heated  to  440°  it  de- 
composes, giving  the  tetroxide  P204  and  red  phosphorus. 

Acids  of  Phosphorus.  —  There  are  six  different  acids  of  phos- 
phorus. Three  are  phosphoric  acids,  representing  the  same  stage 
of  oxidation  of  phosphorus,  but  different  degrees  of  hydration  of 
the  anhydride.  The  others  show  three  different  and  lower  states 
of  oxidation  (study  by  positive  and  negative  valences,  p.  323) : 

Orthophosphoric  acid     H3PO4  ( =  3H2O,P2O5) 
Pyrophosphoric  acid       H4P2O7  ( =  2H2O,P2O6) 
Metaphosphoric  acid      HPO3  (=  H2O,P2O6) 
Hypophosphoric  acid      H2PO3  (=  2H2O,P2O4) 
Phosphorous  acid  H3PO3  ( =  3H2O,P2O3) 

Hypophosphorous  acid  H3PO2  (=  3H2O,P2O) 

The  Phosphoric  Acids.  —  The  relation  between  the  three 
different  phosphoric  acids  may  be  seen  by  considering  them  as 
being  formed  from  phosphorus  pentoxide  (the  anhydride)  and 
water.  In  the  majority  of  cases  already  considered  this  sort  of 
action  takes  place  in  but  one  way.  Thus,  nitric  acid  is  known  in 
but  one  form,  which  is  produced  by  the  union  of  one  molecule 
each  of  nitrogen  pentoxide  and  water :  N2O6  -f  H2O  — >  2HNO3. 
Similarly,  the  chief  sulphuric  acid  is  the  one  formed  from  one 
molecule  of  sulphur  trioxide  and  one  molecule  of  water:  SO3  -+• 
H2O  — >  H2SO4,  although  here  we  have  also  disulphuric  acid 
H2S207,  or  H20,2SO3. 

Now,  when  phosphoric  anhydride  acts  upon  water  we  obtain  a 
solution  which,  on  immediate  evaporation,  leaves  a  glassy  solid, 
HPO3,  known  as  metaphosphoric  acid.  This  is  H2O,P205.  When, 
however,  the  solution  is  allowec}  to  stand  for  some  days,  or  is 
boiled  with  a  little  dilute  nitric  acid,  whose  hydrogen-ion  acts 
catalytically,  the  residue  from  evaporation  is  H3PO4,  orthophos- 
phoric  acid: 

P205  +  3H20^2H3P04    or    HP03  +  H2O  ->  H3PO4. 

This  acid  is  3H2O,P205,  and  no  further  addition  of  water  to  form 
a  different  acid  (see  p.  370)  can  be  effected. 

Conversely,  when  orthophosphoric  acid  is  kept  at  about  255° 
for  a  time,  it  slowly  loses  water,  and  H4P2O7,  pyrophosphoric  acid, 
is  obtained : 


PHOSPHORUS 


369 


This  acid  is  2H2O,P2O5.  Further  desiccation  leaves  metaphos- 
phoric  acid,  which  cannot  be  further  resolved  into  phosphorus 
pentoxide  and  water.  When  dissolved  in  water,  pyrophosphoric 
acid  slowly  resumes  the  water  which  it  has  lost  and  gives  the  ortho- 
acid  again.  The  relations  of  all  these  substances  are  more  clearly 
seen  in  the  graphic  formulae: 


O 

O-H 
0-H 
O-H 

O-H 
O-H 
O-H 

O 


=  0 

=  0 

-O-H 

P 

-o- 

H 

-O-H 

=  0 

=  0 

—  -> 

-O-H 

—  o 

-O-H 

P 

-O- 

H 

P  - 


=  0 

=  0 

O 

=  0 
=  0 


A  most  important  fact  to  be  noted  is  that  the  addition  or  removal  of 
water  leaves  the  valence  of  the  phosphorus  unchanged.  The  degree  of 
oxidation  of  the  phosphorus  and  its  valence  are  identical  in  the 
three  acids. 

The  Relations  of  Anhydrides  and  Oxygen  Acids.  —  Con- 
sidering the  anhydride  from  which  an  acid  is  derived  gives  us  the 
key  to  the  nature  and  behavior  of  the  acid.  It  tells  much  that 
the  formula  of  the  acid  does  not  tell.  For  example:  (1)  What  is 
the  valence  of  phosphorus  in  H3POs?  The  only  way  to  answer 
this  question  is  to  resolve  the  formula  (doubled,  if  necessary)  into 
water  and  the  anhydride,  3H20,P2O3.  The  phosphorus  is  triva- 
lent.  (2)  How  is  this  acid  related  to  metaphosphoric  acid  HP03? 
Resolve  the  latter,  as  before,  H2O,P205.  The  answer  is  that  in  the 
latter  the  phosphorus  is  quinquivalent.  (3)  How  can  we  get 
phosphorous  acid  from  metaphosphoric  acid,  or  vice  versa?  Con- 
sidering the  anhydrides,  we  answer,  by  reduction  and  oxidation, 
respectively.  (4)  Is  pyrophosphoric  acid  H4P2O7,  because  it  con- 
tains 70,  to  be  made  from  all  others  by  oxidation?  Resolve  it  into 
water  and  anhydride,  2H20,P2O5.  We  then  perceive  that  to  make 
it  from  phosphorous  acid  requires  oxidation,  but  to  make  it  from 
ortho-  or  metaphosphoric  acid  requires  only  a  change  in  the  de- 
gree of  hydration:  adding  or  subtracting  water,  since  it  adds  or 
subtracts  hydrogen  and  oxygen  in  equivalent  amounts,  is  not 


370  COLLEGE    CHEMISTRY      , 

.          v 

^oxidation  or  reduction.    These  conceptions  haVe  been  discussed  be- 
fore (pp.  316,  321).  , 

Orthophosphoric  Acid  H3PO^. — The  impure,  commercial  acid 
is  made  by  mixing  selected,  pulverized  phosphate  rock  Ca3(P04)2 
with  sulphuric  acid  (sp.  gr.  1.5)  and  heating  with  steam  and 
stirring  in  a  wooden  vat. 

Ca3(P04)2  +  3H2S04  fc?  2H3P04  +  3CaS04| . 

The  calcium  sulphate  is  precipitated  during  the  heating  and  the 
subsequent  concentration  of  the  filtrate. 

Pure  orthophosphoric  acid  may  be  made  by  boiling  red  phos- 
phorus with  slightly  diluted  nitric  acid  and  evaporating  the  water 
and  excess  of  nitric  acid.  The  product  of  recrystallization  is  a 
white,  crystalline,  deliquescent,  solid  hydrate,  2H3PO4,H20. 

The  acid  is  much  weaker  than  sulphuric  acid,  and  is  dissociated 
chiefly  into  the  ions  H+  and  H2P04~.  The  dihydrophosphate-ion 
H2PO4~,  being^an  acid  as  well  as  an  ion,  is  further  broken  up  to  some 
extent  into  H+  and  HP04=,  as  we  learn  from  the  fact  that  the 
solution  of  the  sodium  salt  NaH2P04  is  acid.  The  ion  HPO4—  is 
hardly  dissociated  at  all,  for  a  solution  of  the  salt  Na2HP04  is 
not  acid  in  reaction. 

Salts  of  Orthophosphoric  Acid.  —  As  a  tribasic  acid,  it  forms 
salts  of  three  kinds,  such  as  NaH2PO4,  NaaHPO4,  and  NasPO4. 
These  are  known  respectively  as  primary,  secondary,  and  tertiary 
sodium  orthophosphate.  The  primary  sodium  phosphate  is 
faintly  acid  in  reaction.  The  secondary  one  is  slightly  alkaline, 
because  of  hydrolysis  arising  from  the  tendency  of  the  hydrogen- 
ion  of  the  water  to  combine  with  the  HP04—  to  form  H2PO4~, 
which  is  much  more  feebly  acid  than  is  phosphoric  acid  H3PO4. 
The  simplified  equation  (p.  271)  shows  the  reason  for  the  alka- 
linity of  the  solution:  HP04=  +  H+  +  OH~  -»  H2P04~  +  OH", 
for  hydroxyl-ion  is  present.  The  tertiary  phosphate  is  stable  only 
in  solid  form,  and  can  be  made  by  evaporating  to  dryness  a 
mixture  of  the  secondary  phosphate  and  sodium  hydroxide : 
Na.HPO,  +  NaOH  <=>  Na3PO4  +  H20 1 . 

When  the  product  is  dissolved  in  water,  this  action  is  reversed  (cf. 
p.  271).  Mixed  phosphates  are  also  known,  particularly  sodium- 


PHOSPHORUS  371 

ammonium  phosphate  (microcosmic  salt)  NaNH4HPO4,4H2O, 
and  the  insoluble  magnesium-ammonium  phosphate  MgNEUPC^. 
Primary  calcium  phosphate  (q.v.),  known  in  commerce  as  "  super- 
phosphate," is  used  as  a  fertilizer. 

The  tertiary  phosphates  are  unchanged  by  heating.  The  primary 
and  secondary  phosphates,  however,  retaining,  as  they  do,  some  of 
the  original  hydrogen  of  the  phosphoric  acid,  are  capable  of  losing 
water,  like  phosphoric  acid  itself,  when  heated.  The  actions  are 
slowly  reversed  when  the  products  are  dissolved  in  water: 

NaH2P04^NaP03    +H20t. 
2Na2HPO4  <=*  Na4P2O7  +  H2O  f . 

It  will  be  seen  that  the  meta-  and  pyrophosphates  of  sodium  are 
formed  by  these  actions;  and  this  is  indeed  the  simplest  way  of 
forming  these  substances,  since  the  acids  themselves  are  not  per- 
manent in  solution,  and  are  too  feeble  to  lend  themselves  to  exact 
neutralization.  Ammonium  salts  of  phosphoric  acid  lose  am- 
monia, as  well  as  water,  when  heated  (cf.  p.  344,  last  par.).  Thus, 
microcosmic  salt  gives  primary  sodium  phosphate: 

NaNH4HP04  ->  NH3 1  +  NaH2P04  -»  NaP03  +  H20 1 , 

and  this  in  turn  is  converted  into  the  metaphosphate  by  loss  of 
water. 

Pyrophosphoric  Acid  and  Metaphosphoric  Acid.  —  Pyro- 
phosphoric  acid  HJ^Oj,  although  tetrabasic,  gives  only  the  neutral 
salts,  such  as  Na4P2O7,  and  those  in  which  one-half  of  the  hydrogen 
has  been  displaced  by  a  metal,  such  as  Na2H2P2O7. 

Metaphosphoric  acid  HP03  is  the  "glacial  phosphoric  acid"  of 
commerce,  and  is  usually  sold  in  the  form  of  transparent  sticks.  It 
is  obtained  by  heating  orthophosphoric  acid,  or  by  direct  union  of 
phosphorus  pentoxide  with  a  small  amount  of  cold  water.  It 
passes  into  vapor  at  a  high  temperature,  and  its  vapor  density 
corresponds  to  the  formula  (HP03)2. 

Sodium  metaphosphate  NaP03,  in  the  form  of  a  small  globule 
obtained  by  heating  microcosmic  salt  on  a  platinum  wire,  is  used 
in  analysis.  When  minute  traces  of  oxides  of  certain  metals  are 
placed  upon  such  a  globule,  known  as  a  bead,  and  heated  in  the 
Bunsen  flame,  the  mass  is  colored  in  various  tints  according  to  the 


372  COLLEGE    CHEMISTRY 

oxide  used  (bead  test).  This  action  may  be  understood  when  we 
consider  that  sodium  metaphosphate  takes  up  water  to  form 
primary  sodium  orthophosphate  :  NaPO3  +  H2O  —  >  NaH2PO4. 
In  the  same  way,  but  at  higher  temperatures,  it  is  able  to  take  up 
oxides  of  elements  other  than  hydrogen,  giving  mixed  ortho- 
phosphates.  Thus,  with  oxide  of  cobalt  a  part  of  the  metaphos- 
phate unites  according  to  the  equation  : 

NaP03  +  CoO  ->  NaCoP04, 
and  the  product  gives  a  blue  color  to  the  bead. 

Distinguishing  Tests.  —  When  a  solution  of  nitrate  of  silver  is 
added  to  a  solution  of  orthophosphoric  acid,  or  to  any  soluble 
orthophosphate,  a  yellow  precipitate  of  silver  orthophosphate 
AgsPO4  is  produced.  This  is  a  test  for  orthophosphate-ion.  With 
pyrophosphoric  acid  or  any  pyrophosphate  the  product  is  white 
Ag4P2O7.  With  metaphosphoric  acid  a  white  precipitate,  AgP03, 
is  obtained  also.  Metaphosphoric  acid  coagulates  a  clear  solu- 
tion (colloidal  suspension)  of  albumin  (say,  white  of  egg),  while 
ortho-  or  pyrophosphoric  acid  has  no  visible  effect  upon  it  (p.  417). 

Phosphorous  Acid  HSPO3.  —  When  added  to  cold  water,  phos- 
phorus trioxide  (P^e)  yields  phosphorous  acid  very  slowly.  With 
hot  water  the  action  is  exceedingly  violent  and  complex  (p.  367). 
This  acid  may  be  obtained  easily  by  the  action  of  water  upon 
phosphorus  trichloride,  tribromide  (p.  197),  or  tri-iodide  and 
evaporation  of  the  solution  : 

3H20  ->  P(OH)3  +  3HC1  1  . 


Some  of  this  acid,  along  with  phosphoric  acid  and  hypophosphoric 
acid,  is  formed  when  moist  phosphorus  oxidizes  in  the  air. 

In  spite  of  the  presence  of  three  hydrogen  atoms,  this  acid  is 
dibasic,  and  two  only  are  replaceable  by  metals.  To  express  this 
fact,  the  first  of  the  following  formulae  is  preferred  : 


H 

0-H 

0-H 


-O-H 
-0-H 
-O-H 


since  the  symmetrical  formula  would  indicate  no  difference  be- 
tween the  three  hydrogen  atoms.     H  united  directly  to  P,  as 


PHOSPHORUS  373 

here  and  in  PH3,  is  not  acidic.  Phosphorous  acid  is  a  powerful 
reducing  agent,  precipitating  silver,  for  example,  in  the  metallic 
form  from  solutions  of  its  salts.  When  heated,  it  decomposes, 
giving  the  most  stable  acid  of  phosphorus  (cf.  pp.  290,  308,  314, 
320,  357),  namely,  metaphosphoric  acid,  and  phosphine: 

4H3P03  ->  3HP03  +  3H20  +  PH3. 

Sulphides  of  Phosphorus. — White  phosphorus,  when  heated 
with  sulphur,  unites  with  explosive  violence.  By  using  red  phos- 
phorus the  action  can  be  controlled.  By  employing  the  proper 
proportions  the  pentasulphide  P2Ss  is  secured.  It  is  purified  by 
distillation,  and  is  a  yellow  crystalline  solid  (m.-p.  274°,  b.-p. 
530°).  Phosphorus  pentasulphide  is  hydrolyzed  by  cold  water: 

P2S5  +  8H20  ->  2H3P04  +  5H2S. 

Other  sulphides,  P4S3  (used  in  making  matches),  P2S3,  and 
P3SC,  may  be  prepared  by  using  the  constituents  in  the  proportions 
represented  by  these  formulae. 

Comparison  of  Phosphorus  with  Nitrogen  and  with  Sul- 
phur. —  Although  phosphorus  and  nitrogen  are  regarded  as  be- 
longing to  one  family,  the  differences  between  them  are  more 
conspicuous  than  the  resemblances.  The  latter  are  confined  al- 
most wholly  to  matters  concerned  with  valence.  The  differences 
are  seen  in  the  facts  that  nitrogen  is  a  gas  and  exists  in  but  one 
form,  while  phosphorus  is  a  solid  occurring  in  two  varieties,  and 
that  the  former  is  inactive  and  the  latter  active.  The  contrasts 
between  phosphine  and  ammonia  (p.  366)  and  between  the  halides 
of  the  two  elements  (pp.  346,  367)  have  been  noted  already. 
The  pentoxide  of  nitrogen  decomposes  spontaneously^  that  of 
phosphorus  is  one  of  the  most  stable  of  compounds.  Nitric  acid 
is  very  active,  both  as  acid  and  oxidizing  agent;  the  phosphoric 
acids  are  quite  the  reverse. 

On  the  other  hand,  the  resemblance  of  phosphorus  to  sulphur 
is  marked.  Both  are  solids,  existing  in  several  forms.  Both 
yield  stable  compounds  with  oxygen  and  chlorine.  The  hydrogen 
compounds  interact  with  salts  to  give  phosphides  of  metals  and 
sulphides  of  metals,  respectively.  Against  these  must  be  set  the 
facts,  that  hydrogen  sulphide  does  not  unite  with  the  hydrogen 


374  COLLEGE    CHEMISTRY 

halides  at  all  while  phosphine  gives  the  phosphonium  halides,  and 
that  phosphoric  acid  is  hard  to  reduce  while  sulphuric  acid  is  re- 
duced with  comparative  ease. 

Exercises.  —  1.  What  are  the  valences  of  the  non-metals  in: 
H2S207,  H2Cr207,  KMn04,  KH2PO2,  H3NO4,  NaH2PO3,  Na^P03? 
Name  these  substances. 

2.  Is  it  oxidation  or  reduction,  or  neither,  when  we  make,  (a) 
N2O4  from  HNO3,  (6)  SO2  from  H2SO3,  (c)  HP03  from  H3P03,  (d) 
H2S207  from  H2S04,  (e)  Na2S04  from  NaHS03? 

3.  Why  would  a  mixture  of  potassium  dichromate  and  hydro- 
chloric acid  (p.  270)  be  less  suitable  than  nitric  acid,  as  an  oxidizing 
agent  for  making  phosphoric  acid  irom  red  phosphorus? 

4.  Why  is  not  the  tertiary  phosphate  of  sodium  (p.  371)  decom- 
posed by  heating?    \Vhat  tertiary  phosphates  would  be  decom- 
posed by  this  means? 

5.  Formulate  the  hydrolyses  of  the  secondary  and  tertiary 
sodium  orthophosphates  as  was  done  for  sodium  sulphide  (p.  271). 

6.  How  should  you  prepare  Ca2P207  and  Ca(P03)2,  both  in- 
soluble? 

7.  What  product  should  you  confidently  expect  to  find  after 
heating  (p.  371),  (a)  sodium  phosphite,  Na^HPOs,  (6)  potassium 
hypophosphite?     Make  the  equations. 

8.  Compare  the  elements  chlorine  and  phosphorus  after  the 
manner  of  the  comparisons  on  p.  373. 


CHAPTER  XX1 
CARBON  AND   THE   OXIDES   OF   CARBON 

THE  majority  of  the  substances  composing,  or  produced  by, 
living  organisms,  such  as  starch,  fat,  and  sugar,  are  compounds  of 
carbon.  Hence  the  chemistry  of  these  compounds  is  known  as 
organic  chemistry.  It  was  at  first  supposed  that  the  artificial 
production  of  such  compounds,  e.g.,  without  the  intervention  of  life, 
was  impossible.  But  many  natural  organic  products  have  now 
been  made  from  simpler  ones  or  from  the  elements,  and  the  prepa- 
ration of  the  others  is  delayed  only  in  consequence  of  difficulties 
caused  by  their  instability  and  complexity.  On  the  other  hand, 
hundreds  of  compounds  unknown  to  animal  or  vegetable  life,  in- 
cluding many  valuable  drugs  and  dyes,  have  now  been  added  to 
the  catalogue  of  chemical  compounds.  More  than  200,000  differ- 
ent compounds  containing  carbon  are  known,  and  thousands  are 
added  every  year. 

The  elements  entering  into  carbon  compounds  are  chiefly  hydro- 
gen and  oxygen.  After  these,  nitrogen,  phosphorus,  the  halogens 
and  sulphur  may  be  named. 

CARBON  C 

Occurrence.  —  Large  quantities  of  carbon  are  found  in  the  free 
condition  in  nature.  The  diamond  is  the  purest  natural  carbon. 
Graphite,  or  plumbago,  which  is  the  next  purest,  is  found  in 
limited  amounts,  and  is  a  valuable  mineral.  Coal  occurs  in 
numerous  forms  containing  greatly  varying  proportions  of  free 
carbon.  Small  quantities  of  the  free  element  have  been  found  in 
meteorites. 

In  combination,  carbon  is  found  in  marsh-gas,  or  methane  CHi, 
which  is  the  chief  component  of  natural  gas.  The  numerous  com- 
pounds found  in  plants  and  animals  have  already  been  mentioned. 
The  mineral  oils  consist  almost  entirely  of  mixtures  of  various 
compounds  of  carbon  and  hydrogen.  Whole  geological  formations 

375 


376  COLLEGE    CHEMISTRY 

are  composed  of  carbonates  of  common  metals,  particularly  calcium 
carbonate  or  limestone. 

Allotropic  Forms  of  Carbon.  —  The  allotropic  (p.  222)  forms 
of  carbon  differ  very  strikingly  in  their  physical  properties.  The 
diamond  is  transparent,  crystalline,  and  very  hard  (sp.  gr.  3.5). 
Graphite  is  black,  lustrous,  and  very  soft  (sp.  gr.  2.3).  Amorphous 
carbon  is  very  variable.  Thus  lampblack  (see  p.  398)  is  a  fine 
powder  of  nearly  pure  carbon,  charcoal  (see  p.  408)  shows  the 
structure  of  the  wood,  and  coal  (see  p.  409)  contains  compounds  of 
carbon  as  well  as  the  free  element.  These  amorphous  forms  can 
best  be  discussed  after  the  materials  from  which  they  are  formed 
have  been  considered. 

That  all  the  forms  are  composed  of  the  same  element  is  shown 
by  the  fact  that  they  all  burn  in  oxygen  to  give  carbon  dioxide. 
Then,  too,  when  heated  strongly  in  absence  of  air,  diamond  and 
the  amorphous  forms  all  turn  into  graphite.  They  contain  differ- 
ent amounts  of  chemical  energy,  however.  Thus,  when  1  g.  of 
each  is  burned,  diamond  gives  7805  cal.,  graphite  7850  and  sugar 
charcoal  (p.  286)  8040.  The  tendency  of  most  carbon  compounds, 
when  heated,  to  char,  giving  free  carbon,  is  used  as  a  test. 

The  Diamond.  —  Diamonds  are  found  chiefly  in  Brazil,  and 
South  Africa.  They  are  separated  by  weathering  the  rock,  which 

then  crumbles,  and  by  washing  the 
debris  with  water.  They  are  covered 
with  a  crust  which  entirely  obscures 
their  luster,  and  possess  natural  crys- 
talline forms  belonging  to  the  regu- 
lar system,  such  as  the  octahedron 
(p.  83).  It  should  be  noted  that 
FlG  99  this  natural  form  bears  no  relation 

whatever  to  the  pseudo-crystalline 

shape  which  is  conferred  upon  the  stone  by  the  diamond-cutter. 
The  natural  stone  is  "cut,"  by  grinding  new  faces.  Thus,  a 
"brilliant"  possesses  one  rather  large,  flat  face,  which  forms  the 
base  of  a  many  sided  pyramid  (Fig.  99,  showing  two  views) .  This 
form  is  given  to  the  stone,  in  order  that  the  maximum  reflection 
of  light  from  its  interior  may  be  produced.  The  diamond  is  harder 


CARBON  AND  THE  OXIDES  OF  CARBON 


377 


(Appendix  II)  than  any  other  variety  of  matter,  so  that  it  can  be 
scratched  or  polished  only  by  rubbing  with  diamond  powder.  It 
is  the  densest  form  of  carbon  (sp.  gr.  3.5).  The  colorless  stones, 
and  occasional  specimens  with  special  tints  (like  the  blue,  Hope 
diamond)  are  the  most  valuable.  The  black  ("carbonado")  and 
discolored  specimens  are  used  for  grinding  and  glass  cutting. 
Mounted  round  the  edge  of  a  tube,  they  are  used  for  drilling  rock, 
so  that  a  cylindrical  specimen  of  the  whole  of  the  strata  can  be 
secured  for  examination.  The  forms  of  carbon  are  insoluble  in  all 
liquids  at  room  temperature.  Molten  iron  (q.v.)  dissolves  five  or 
six  per  cent,  part  of  which  goes  into  combination;  but  usually  only 
graphite  is  found  in  the  cooled  product.  Moissan  (1887),  however, 
succeeded  in  preparing  microscopic  fragments  of  diamonds  in  this 
way.  The  diamond  is  a  nonconductor  of  electricity. 

Diamonds  are  sold  by  the  new  international  carat,  200  mgms. 
(old  carat,  4  grains  =  205  mgms.),  and  the  value  increases  with 
the  size.  Thus,  a  first  quality,  cut  stone  of  1  carat  is  worth  about 
$270,  one  of  2  carats  about  $340  per  carat.  The  largest  diamond 
known,  the  Cullinan  (1905),  weighed  3032  (old)  carats  (621  g.  or 
1.37  Ibs).  It  was  presented  by  the  Transvaal  government  to 
King  Edward  VII,  and  was  cut  into  stones  of  516.5  and  309 
carats  and  many  smaller  ones.  Other  large  stones  are  the  Jubilee 
(239  carats),  and  the  Kohinoor  (106  carats). 

Graphite.  —  Graphite  (Gk.,  I  write)  is  found  in  Cumberland, 
Siberia,  Canada,  and  Ceylon.  It  is  composed  of  glittering,  slip- 
pery, crystalline  scales  (hexa- 
gonal system).  In  utter  con- 
trast to  the  diamond,  the  min- 
eral is  extremely  soft,  has  a 
smaller  specific  gravity  (2.3), 
and  conducts  electricity.  It  is 
made  artificially  by  an  electro- 
thermal process  (cf.  p.  363).  A 
powerful  alternating  current  is 

passed  through  a  mass  of  granular  anthracite,  mixed  with  pitch  and 
a  little  sand  (Acheson's  process).  The  mixture  (3  tons)  is  piled 
between  the  electrodes  (Fig.  100)  and,  on  account  of  its  high  resist- 
ance, becomes  strongly  heated.  The  change  occupies  24-30  hours. 


378  COLLEGE    CHEMISTRY 

Graphite  is  now  used  exclusively  for  making  the  anodes  in  the 
electrolytic  manufacture  of  chlorine  and  in  related  processes. 
Mixed  with  fine  clay  it  forms  the  "lead"  of  lead  pencils,  first  used 
in  the  sixteenth  century.*  Mixed  with  clay  it  is  used  also  for 
making  crucibles,  which  withstand  high  temperatures  and  serve 
for  melting  and  casting  steel  and  high  melting  alloys.  As  "black- 
lead"  it  forms  stove  polish,  the  layer  of  fine  scales  protecting  the 
iron  against  rusting.  It  is  employed  as  a  lubricant  in  cases  whore 
oil  would  be  decomposed  by  the  heat  and  where  wooden  surfaces 
are  in  contact. 

Chemical  Properties  of  Carbon.  —  The  most  common  uses 
of  carbon  depend  upon  its  great  tendency  to  unite  with  oxygen, 
forming  carbon  dioxide  CO2.  Under  some  circumstances  carbon 
monoxide  CO  (see  below)  is  produced.  Aside  from  the  direct 
employment  of  this  action  for  the  sake  of  the  heat  which  is  liber- 
ated, it  is  used  also  in  the  reduction  of  ores  of  iron,  copper,  zinc,  and 
many  other  metals.  When,  for  example,  finely  powdered  cupric 
oxide  and  carbon  are  heated,  copper  is  obtained.  The  gas  given 
off  is  either  carbon  dioxide,  or  a  mixture  of  this  with  carbon  mon- 
oxide, according  to  the  proportion  of  carbon  used: 

CuO  +  C->   Cu  +  CO, 


At  the  high  temperatures  produced  in  the  electric  furnace,  carbon 
unites  with  many  metals  and  some  non-metals.  Compounds 
formed  in  this  way  are  known  as  carbides,  such  as  aluminium 
carbide  AUCs,  calcium  carbide  CaC2,  and  carborundum  CSi  (see 
below)  . 

The  union  with  hydrogen  is  ordinarily  too  slow  to  be  observed. 
But  when  the  carbon  is  mixed  with  pulverized  nickel  (contact 
agent),  and  hydrogen  is  passed  over  the  mixture  at  250°,  methane 
CH4  is  formed  (99  per  cent).  The  action  is  reversible  and  exo- 
thermal, and  is  therefore,  at  higher  temperature,  less  complete 
(cf.  p.  189),  at  850°  reaching  only  1.5  per  cent.  On  the  other  hand, 
an  electric  arc,  between  carbon  poles  in  an  atmosphere  of  hydrogen, 
gives  traces  of  acetylene  C2H2,  this  action  being  endothermal.  The 

*  Priestley  was  the  first  to  suggest  the  use  of  caoutchouc  (raw  rubber)  as 
an  eraser. 


CARBON   AND    THE    OXIDES    OF   CARBON  379 

other  compounds  of  carbon  and  hydrogen  are  all  obtained  by  in- 
direct reactions. 

Carbon  Disulphide  CSQ.  —  This  compound  is  made  by  direct 
union  of  sulphur  vapor  and  glowing  charcoal.  An  electric  furnace 
like  that  in  Fig.  98  (p.  363)  is  employed.  The  substance  comes  off 
as  a  vapor  and  is  condensed. 

Carbon  disulphide  is  a  colorless,  highly  refracting  liquid  (b.-p. 
46°).  Traces  of  other  compounds  give  the  commercial  article  a 
disagreeable  smell.  It  burns  in  air,  forming  carbon  dioxide  and 
sulphur  dioxide.  It  is  an  important  solvent  for  sulphur  and 
caoutchouc  (rubber),  and  dissolves  iodine  and  phosphorus  freely. 
Large  quantities  are  employed  also  in  the  destruction  of  prairie 
dogs  and  ants  and  for  freeing  grain  elevators  of  rats  and  mice. 

Carbon  Tetrachloride  CC/4-  —  This  compound  is  manu- 
factured by  leading  dry  chlorine  into  carbon  disulphide  containing 
a  little  iodine  (contact  agent)  in  solution: 


The  carbon  tetrachloride  (b.-p.  77°)  is  first  distilled  off,  and  the 
sulphur  monochloride  (b.-p.  136°)  is  purified  for  use  in  vulcanizing 
rubber. 

Carbon  tetrachloride  is  a  colorless  liquid  which  dissolves  fats, 
fcars,  and  many  other  organic  compounds.  It  is  used  to  take  the 
oil  or  grease  out  of  wool,  linen,  oil-bearing  seeds  and  bones.  It  has 
the  advantage  over  gasoline  (petrol)  and  benzine  (see  p.  391), 
which  can  be  used  for  similar  purposes,  that  it  is  non-inflammable. 
"Carbona,"  used  for  removing  stains  from  clothing,  gloves  and 
shoes,  is  benzine  to  which  sufficient  carbon  tetrachloride  has-been 
added  to  render  the  mixture  noninflammable.  "Pyrene" 
extinguishers  contain,  mainly,  carbon  tetrachloride.  The  tem- 
perature of  the  burning  material  is  lowered,  because  heat  is  con- 
sumed in  vaporizing  the  liquid,  and,  at  the  same  time,  the  vapo: 
displaces  the  air  and  stops  the  combustion. 

Calcium  Carbide  CaC2  and  Carborundum  SiC.  —Calcium 
carbide  is  manufactured  in  an  electric  furnace,  by  the  interact: 
of  finely  pulverized  limestone  or  quicklime  with  coke: 

3C->CaC2  +  CO. 


380 


COLLEGE    CHEMISTRY 


The  operation  is  a  continuous  one,  the  materials  being  thrown  into 
the  left  side  of  the  drum  (Fig.  101,  diagrammatic),  and  the  product 
removed  on  the  right.  The  carbon  poles  are  fixed.  The  arc  having 
been  established,  the  drum  is  rotated  slowly  as  the  carbide  accu- 
mulates. The  current  enters  by  one  carbon,  passes  through  the 
carbide,  and  leaves  by  the  other.  The  high  resistance  of  the 
partially  transformed  material  causes  the  production  of  the  heat. 
When  the  action  in  one  layer  approaches  completion,  the  resistance 
falls,  the  current  increases,  and  an  armature  round  which  the  wire 
passes  (not  shown  in  Fig.  101)  comes  into  operation  and  turns  the 
drum.  In  this  way  the  carbide  just  formed,  is  continuously  moved 

away  from  the  carbons,  and 
new  material,  introduced  on 
the  left,  falls  into  the  path 
of  the  current.  The  iron 
plates  which  form  the  cir- 
cumference of  the  drum  are 
added  on  the  left  and  re- 
moved on  the  right,  where 
also  the  carbide  is  broken 
out  with  a  chisel.  The  drum 
revolves  once  in  about  three 
days.  The  product  is  used 
for  making  acetylene  (q.v.). 

Carborundum,  or  carbide 
of  silicon  SiC,  of  which 
hundreds  of  tons  are  manu- 
factured annually  at  Niagara  Falls  (Acheson's  process),  is  made 
in  an  electric  furnace  of  the  type  shown  in  Fig.  100  (p.  377).  A 
mixture  of  coke  and  sand  (silicon  dioxide  Si02)  with  some  saw- 
dust is  piled  between  the  terminals,  with  a  core  of  granular  carbon 
to  carry  most  of  the  current.  The  resistance  produces  a  high 
temperature  (1950°),  and  carbon  replaces  the  oxygen: 

3C  +  Si02-»SiC  +  2CO. 

The  carborundum  remains,  often  in  beautifully  crystalline  form. 
It  is  exceedingly  hard  (Appendix  II),  and  after  pulverization  and 
mixing  with  a  filler,  is  moulded  into  grinding  wheels. 


Fia.  101. 


CARBON  AND   THE   OXIDES   OF   CARBON  381 

CARBON  DIOXIDE  AND  CARBONIC  ACID 

Occurrence.  —  Carbon  dioxide  is  present  in  the  atmosphere, 
and  issues  from  the  ground  in  large  quantities  in  certain  neighbor- 
hoods, as,  for  example,  in  the  so-called  Valley  of  Death  in  Java, 
and  in  the  Grotta  del  Cane  near  Naples.  Effervescent  mineral 
waters,  such  as  those  of  Vichy  and  of  the  Geyser  Spring  at  Sara- 
toga, contain  it  in  solution,  and  their  effervescence  is  caused  by  the 
escape  of  the  gas  when  the  pressure  is  reduced. 

Modes  of  Formation.  —  1.  Carbon  dioxide  is  produced  by 
combustion  of  carbon  with  an  excess  of  oxygen:  C  +  02  —  »  C02. 
The  combustion  of  all  compounds  of  carbon,  as  well  as  the  slow 
oxidation  in  the  tissues  of  plants  and  animals,  yield  the  same 
product.  The  product  from  burning  carbon  is  naturally  mixed 
with  at  least  four  times  its  volume  of  atmospheric  nitrogen.  To 
secure  carbon  dioxide  for  commercial  purposes  from  this  source, 
the  gas  is  led  under  pressure  into  a  solution  of  potassium  carbonate, 
which  absorbs  the  carbon  dioxide: 

C02  (gas)  <=>  C02  (dslvd)  +  H20  <=>  H2C03  +  K2C03  <=t  2KHC03. 

When  the  pressure  is  reduced  by  a  pump,  all  the  actions  are 
reversed,  and  the  gas  escapes  in  pure  form.  The  same  solution, 
with  occasional  purification,  can  be  used  an  indefinite  number  of 
times. 

2.  It  was  Joseph  Black  (1757)  who  first  recognized  the  gas  as  a 
distinct  substance.  He  observed  its  formation  when  marble  or 
magnesium  carbonate  was  heated: 

CaCO3<=±CaO 


and  named  the  gas  "fixed  air"  from  the  fact  that  it  was  contained 
in  these  solids.  The  above  action  had  been  used  for  centuries 
in  making  quicklime  (calcium  oxide).  All  common  carbonates, 
excepting  the  normal  carbonates  of  potassium  and  sodium,  de- 
compose, leaving  the  oxide  of  the  metal  or  the  metal  (p.  60). 

3.  Black  found  that  the  gas  was  also  produced  when  acids  acted 
upon  carbonates,  and  this  is  the  method  employed  in  the  labora- 
tory: 

CaC03  (solid)  ±5  CaC03  (dslvd)  fc?  Ca^  +  CO3=  1       R  m  «-  H  O  4-  CO. 
2HC1  (dslvd)  1=5201-  +2H+    f- 


382  COLLEGE    CHEMISTRY 

Since  the  carbonic  acid  is  very  slightly  ionized,  the  action  is  like 
that  of  acids  on  sulphites  (p.  275).  Since,  however,  the  carbonate 
of  calcium  (marble)  is  very  slightly  soluble,  so  that  an  additional 
equilibrium  controls  its  solution,  the  action  is  like  that  of  acids  on 
ferrous  sulphide  (p.  272).  The  apparatus  shown  in  Fig.  24  is  used. 
Carbon  dioxide  is  formed  in  decay  (p.  36)  and,  as  Black  likewise 
discovered,  in  fermentation  (q.v.). 

Physical  Properties.  —  Carbon  dioxide  is  a  colorless,  odorless 
gas.  It  is  heavier  than  air.  The  G.M.V.  weighs  44.26  g.  Its 
critical  temperature  is  31.35°  (p.  79).  The  solid  melts  at  -56°, 
having  a  vapor  pressure  of  5.3  atmospheres.  The  solid  has  a  vapor 
pressure  of  1  atmos.  at  —79°.  The  sp.  gr.  of  the  liquid  at  0°  is 
0.95.  At  0°  its  vapor  tension  is  35.4  atmospheres  and  at  20°, 
59  atmospheres.  It  must  be  preserved,  therefore,  in  very  strong 
cylinders  of  mild  steel.  Large  quantities  of  it,  often  collected  from 
fermentation  vats,  are  sold  in  such  cylinders,  and  used  in  operating 
beer-pumps  and  in  making  aerated  waters.  When  the  liquid  is 
allowed  to  flow  out  into  an  open  vessel  or,  still  better,  into  a  cloth 
bag  (non-conductor  of  heat),  it  cools  itself  by  its  own  evaporation 
and  forms  a  white,  snowlike  mass.  Solid  carbon  dioxide  evapo- 
rates at  —79°,  without  melting,  since  at  that  temperature  it 
exercises  1  atmosphere  pressure,  and  the  heat  from  the  surround- 
ings is  used  as  heat  of  vaporization  instead  of  being  employed  in 
raising  the  temperature  to  the  melting-point  ( —  56°) . 

The  solid  is  used  in  the  laboratory  as  a  cooling  agent,  being  often 
mixed  with  ether  to  give  closer  contact  with  the  vessel.  Mercury 
(m.-p.  —40°)  is  easily  frozen  by  the  mixture. 

Carbon  dioxide  gas  (760  mm.  and  15°)  dissolves  in  its  own 
volume  of  water.  Up  to  four  or  five  atmospheres  Henry's  law 
(p.  128)  describes  its  solubility  accurately.  An  aqueous  solution, 
under  a  pressure  of  3^4  atmospheres,  is  familiarly  known  as  soda 
water,  or  carbonated  water. 

Chemical  Properties.  —  Carbon  dioxide  is  a  stable  compound. 
At  2000°,  the  dissociation  reaches  1.8  per  cent,  or  about  the  same 
as  that  of  water:  2C02  <=±  2CO  +  O2. 

The  more  active  metals,  like  magnesium,  burn  brilliantly  when 
ignited  in  a  hollow  lump  of  solid  carbon  dioxide,  producing  the 


CARBON  AND  THE    OXIDES   OF   CARBON  383 

oxide  and  free  carbon.  Less  active  metals,  such  as  zinc  and  iron, 
when  heated  in  a  stream  of  the  gas,  give  an  oxide  of  the  metal  and 
carbon  monoxide  (q.v.). 

Carbon  dioxide  unites  directly  with  many  oxides,  particularly 
those  of  the  more  active  metals,  such  as  the  oxides  of  potassium, 
sodium,  calcium,  etc.  Hence  the  decomposition  of  calcium  car- 
bonate by  heating  (p.  381)  is  a  reversible  action. 

Carbon  dioxide,  when  dissolved  in  water,  forms  an  unstable  acid: 

H  O        H-0. 

or  o  +  CT      -»  C  =  O. 


The  name  carbonic  acid  is  frequently,  though  improperly,  given  to 
the  anhydride  CO2,  which  has  no  acid  properties. 

Chemical  Properties  of  Carbonic  Acid  jFr2CO3.  —  The  solu- 
tion of  carbon  dioxide  in  water  exhibits  the  properties  of  a  weak 
acid.  It  conducts  electricity,  although  not  well.  It  turns  litmus 
red.  The  ionization  takes  place  chiefly  according  to  the  equation: 


Carbonates  and  Bicarbonates.  —  When  excess  of  an  aqueous 
solution  of  carbonic  acid  is  mixed  with  a  solution  of  a  base  like 
sodium  hydroxide,  or,  as  the  operation  is  more  usually  performed, 
when  carbon  dioxide  is  passed  into  a  solution  of  the  alkali,  until  the 
liquid  is  saturated,  water  is  formed  and  the  acid  carbonate  (bi- 
carbonate) of  sodium  remains  dissolved: 

H2C03  +  NaOH^H20  +  NaHC03,     or    H+  +  OH~->H20. 

Although  the  bicarbonate  is  technically  an  acid  salt,  its  solution  is 
neutral,  on  account  of  the  exceedingly  slight  dissociation  of  the 
HCO3~  ion.  By  addition  of  an  equivalent  of  sodium  hydroxide, 
the  normal  carbonate  is  obtained: 

NaOH+NaHCO3^H20+Na2C03,  or  OH"+HC03"<=±H20+C03=. 

This  solution,  like  that  of  all  salts  of  a  strong  base  and  a  feeble  acid 
(cf.  p.  271),  is  alkaline  in  reaction.  This  is  because  the  tendency  to 
form  the  very  slightly  ionized  HC03~  makes  the  foregoing  ionic 
action  noticeably  reversible  (cf.  pp.  271,  370). 


384  COLLEGE    CHEMISTRY 

The  normal  carbonates,  with  the  exception  of  those  of  potassium, 
sodium,  and  ammonium,  are  insoluble  in  water,  and  may  be  ob- 
tained by  precipitation  when  the  proper  ions  are  employed.  For 
example  : 

BaCl2+Na«COj<=*BaCOs,l+2NaCl,    or    Ba^+CO3=<=±BaC03j. 

The  aqueous  solution  of  carbon  dioxide  interacts  with  solutions 
of  barium  and  calcium  hydroxides  in  a  similar  manner: 

Ca(OH)2  +  H2CO3  *=+  CaC03|  +  H20. 

These  precipitations  are  used  as  tests  for  carbon  dioxide. 

Excess  of  carbon  dioxide  converts  calcium  carbonate  into  the 
more  soluble  bicarbonate,  ano>  hence  considerable  quantities  of 
"lime"  (hardness,  q.v.)  are  frequently  held  in  solution  by  natural 
waters  which  contain  carbon  dioxide  in  solution  : 

H2C03  +  CaC03  <=>  Ca(HC03)2. 

In  the  same  fashion,  the  carbonates  of  iron  (FeC03),  magnesium, 
and  zinc  are  somewhat  soluble  in  water  containing  free  carbonic 
acid.  In  fact,  the  solution,  transportation,  and  deposition  of  all 
these  carbonates  take  place  in  nature  on  a  large  scale  by  the  alter- 
nate progress  and  reversal  of  this  action. 

Uses  of  Carbon  Dioxide.  —  The  use  of  the  gas  for  impregnat- 
ing aerated  waters  has  been  mentioned.  The  gas  is  used  in  im- 
mense quantities  in  the  manufacture  of  sodium  bicarbonate 
NaHCO3  (baking  soda),  of  sodium  carbonate  Na2CO3,10H2O 
(washing  soda),  and  of  white  lead,  a  basic  carbonate  of  lead 
Pb3(OH)2(C03)2.  > 

Since  carbon  dioxide  is  already  fully  oxidized,  it  does  not  burn, 
and  since  it  is  very  stable,  ordinary  combustibles  will  not  burn  in 
it.  A  small  percentage  of  it  will  destroy  the  power  of  air  to  sup- 
port combustion.  For  this  reason,  portable  fire  extinguishers 
contain  a  dilute  solution  of  sodium  bicarbonate,  and  a  bottle  of 
sulphuric  acid.  When  the  tank  is  inverted,  the  acid  flows  into  the 
solution: 


2NaHC03  +  H2S04  <=±  Na^  +  2H2CO3  <=>  2H2O  +  2C02. 

The  liquid  is  saturated  with  the  gas  and  the  excess,  rising  to  the 
top,  by  its  pressure  forces  the  solution  out  through  the  nozzle.     The 


CARBON   AND   THE    OXIDES   OF   CARBON  385 

liquid  is  more  effective  than  an  equal  amount  of  water,  because 
the  carbon  dioxide  it  carries  mixes  with  the  surrounding  air. 

The  most  wonderful  chemical  change  which  carbon  dioxide 
undergoes  is  perhaps  the  most  useful  to  mankind,  and  at  the  same 
time  the  one  least  understood.  This  is  the  action  by  which  plants 
use  it  as  food  (see  last  section  of  this  chapter). 

CARBON  MONOXIDE  CO 

Preparation.  —  In  the  laboratory,  carbon  monoxide  is  ob- 
tained by  heating  oxalic  acid,  a  solid,  white,  crystalline  substance, 
in  a  flask  with  concentrated  sulphuric  acid.  The  latter  is  here 
employed  simply  as  a  dehydrating  agent  (p.  286),  so  that  it  need 
not  be  included  in  the  equation: 

H2C2O4  -» C02  +  CO  +  H20. 

To  obtain  pure  carbon  monoxide  from  this  mixture,  it  is  necessary 
to  remove  the  carbon  dioxide,  by  passing  the  gas  through  a  solu- 
tion of  potassium  hydroxide  contained  in  a  wash  bottle.  By  using 
formic  acid,  or  sodium  formate,  with  sulphuric  acid,  the  presence 
of  the  carbon  dioxide  is  avoided: 

HCHO2-*CO  +  H20. 

We  commonly  observe  the  blue  flame  of  burning  carbon  mon- 
oxide playing  on  the  surface  of  a  coal  fire.  The  gas  is  produced  by 
the  passage  of  the  carbon  dioxide,  which  is  first  formed,  through 
the  upper  layers  of  heated  coal: 

C-»2CO. 

A  similar  reduction  of  carbon  dioxide  is  produced  when  the  gas  is 
led  over  a  metal,  such  as  zinc,  and  heat  is  applied: 

C02  +  Zn  ->  ZnO  +  CO. 

Producer  Gas  and  Water  Gas.  —  When  coke  and  air  are 
used  in  the  reaction  mentioned  above,  the  mixture  of  carbon 
monoxide  (about  33  per  cent)  and  nitrogen  (about  66  per  cent) 
obtained  is  called  producer  gas.  It  is  combustible  and  is  used  in 
factories  for  heating  and  to  drive  gas  engines  for  power. 


386  COLLEGE    CHEMISTRY 

When  steam  is  driven  through  white  hot  coke  or  anthracite,  a 
mixture  of  hydrogen  and  carbon  monoxide,  known  as  water  gas, 
is  produced- 

C  +  H2O  ->  CO  +  H2  -  28,300  cal. 

The  coke,  piled  in  a  brick-lined,  cylindrical  structure,  is  brought 
to  vigorous  combustion  by  blowing  in  air  for  ten  minutes.  Then 
steam  is  substituted  for  the  air.  Since  the  interaction  takes  place 
with  absorption  of  heat  (is  endothermal,  see  equation),  in  about 
five  minutes  the  coke  "becomes  too  cool.  Air  is  then  substituted 
for  steam,  and  so  on  alternately.  The  gas  is  collected  while  the 
steam  is  turned  on,  and  contains  equal  volumes  of  the  two  gases, 
together  with  some  carbon  dioxide  (4-7  per  cent),  nitrogen  (4-5 
per  cent)  and  oxygen  (1  per  cent).  The  gas  is,  therefore,  almost 
wholly  combustible  and  is  used  as  a  source  of  heat,  and  for  driving 
gas  engines  to  furnish  power.  It  is  used  also  for  making  illuminat- 
ing gas  (q.v.).  Since  carbon  monoxide  is  more  easily  liquefied  than 
is  hydrogen,  the  latter  gas  is  obtained,  for  commercial  use,  by 
passing  water  gas  through  a  liquefier. 

When  both  steam  and  air  are  driven  together  over  burning  coke, 
the  latter  is  able  to  burn  continuously,  and  a  fuel  gas  which  is  a 
cross  between  producer  gas  and  water  gas  is  obtained. 

Fuel  gases  are  used  on  a  large  scale  in  steel  works,  and  other 
factories.  They  give  a  uniform  and  easily  regulated  heat,  they 
leave  no  ash,  and  their  use  involves  no  labor  for  stoking.  As 
gases,  also,  they  can  be  used  in  structures  in  which  coal,  as  a  solid, 
could  not  be  employed. 

Physical  Properties.  —  Carbon  monoxide  is  a  colorless  gas, 
with  a  metallic  odor  and  taste  (poisonous!).  It  is  very  slightly 
soluble  in  water.  Its  density  is  almost  the  same  as  that  of  air. 
When  liquefied  it  boils  at  - 190°. 

Chemical  Properties.  —  All  the  chemical  properties  of  carbon 
monoxide  are  referable  to  the  fact  that  in  it  the  element  carbon 
appears  to  be  bivalent:  CuO.  The  compound  is  in  fact  unsatu- 
rated,  and  combines  with  oxygen,  chlorine,  and  other  substances 
directly.  Thus  the  gas  burns  in  the  air,  uniting  with  oxygen  to 
form  carbon  dioxide.  Again,  iron  (q.v.)  is  manufactured  by  the 


CARBON   AND   THE    OXIDES   OF   CARBON  387 

reduction  of  the  oxide  of  iron  by  gaseous  carbon  monoxide  in  the 
blast  furnace: 

Fe2O3  +.  3CO  <=±  2Fe  +  3CO2. 

In  sunlight  carbon  monoxide  unites  directly  with  chlorine  to  form 
carbonyl  chloride  (phosgene)  COC12.  It  unites  with  nickel  and  iron 
to  form  nickel  carbonyl  and  iron  carbonyl  (q.v.),  respectively. 

The  gas  is  an  active  poison.  When  inhaled  it  unites  with  the 
haemoglobin  of  the  blood,  to  the  exclusion  of  the  oxygen  which 
forms  a  less  stable  compound  (cf.  p.  36).  A  quantity  equivalent 
to  about  10  c.c.  of  the  gas  per  kilo,  weight  of  the  animal  is  sufficient 
to  produce  death,  about  one-third  of  the  whole  haemoglobin  having 
entered  permanently  into  combination  with  carbon  monoxide. 
One  volume  in  800  volumes  of  air  produces  death  in  about  thirty 
minutes.  This  gas  is  the  chief  poisonous  substance  in  illuminating 
gas.  The  poisonous  effect  of  tobacco  smoke,  when  inhaled,  is 
partly  due  to  the  carbon  monoxide  produced  by  incomplete  com- 
bustion. Nicotine,  although  contained  in  tobacco  leaves,  is 
unstable,  and  is  decomposed  by  the  heat.  Traces  of  other  irritant  *' 
organic  compounds,  however,  are  contained  in  the  smoke. 

Carbon  Dioxide  as  Plant  Food.  —  The  walls  of  the  cells 
which  form  the  framework  of  a  plant  are  made  of  cellulose 
(CeHioOs)*.  In  the  cells,  especially  those  in  certain  parts  of  the 
plant,  granules  of  starch  (CeHioOs)^  are  found.  The  plant  con- 
tains also  proteins.  These  substances  contain  carbon,  hydrogen, 
oxygen,  nitrogen,  sulphur,  and  phosphorus,  and  plant  food  must  4 
furnish  these  elements.  Compounds  of  potassium  are  also  re-  i 
quired.  Hence,  in  addition  to  large  amounts  of  water  ascending; 
through  the  roots  and  stem,  carrying  sufficient  quantities  of  solu- 
ble compounds  of  the  four  elements  last  named,  all  plants  require  , 
an  abundant  supply  of  carbon  in  absorbable  form.  Now,  this 
comes  from  atmospheric  carbon  dioxide,  admitted  through  minute 
openings  situated  mainly  on  the  surfaces  of  the  leaves.  Com- 
parison of  the  formulae  CO2  and  C6Hi0O5  shows  at  once  that  the 
assimilation  of  the  carbon  dioxide  of  the  plant  must  involve 
reduction.  The  chlorophyll  (green  matter)  and  protoplasm  in  the 
leaves  act  upon  the  carbon  dioxide,  causing  oxygen  gas  to  be 
liberated : 

6CO2  +  5H2O  +  671,000  cal.  ->  C6Hi0O5  +  6O2. 


388  COLLEGE    CHEMISTRY 

This  action  goes  on  only  in  sunlight,  and  if  green  leaves  are  placed 
under  water  saturated  with  carbon  dioxide,  oxygen  is  given  off 
and  can  be  collected. 

The  enormous  amount  of  energy  absorbed  in  the  action,  and 
represented  in  terms  of  heat  in  the  equation,  is  furnished  by  the 
sunlight.  It  may  be  added  that  plants,  like  animals,  also  use 
some  oxygen  and  produce  some  carbon  dioxide,  but  this  process  is 
entirely  overborne  in  daylight,  and  is  noticeable  only  in  the  dark. 

The  energy  that  does  the  world's  work  comes  mainly  from  two 
sources,  namely,  water  power  and  the  combustion  of  wood  or  coal 
(which  is  fossil  wood).  The  water  comes  from  vapor,  generated 
by  the  sun's  heat,  condensed  as  rain,  and  ultimately  feeding  the 
rivers.  The  source  of  the  energy  in  wood  and  coal  is  now  apparent. 
When  wood,  which  is  largely  cellulose  (CeHioOs)*,  burns,  it  gives 
carbon  dioxide,  water,  and  heat.  In  fact,  its  combustion  is 
represented  by  the  above  equation,  when  read  backwards.  Thus, 
the  sunlight,  through  the  machinery  of  the  plant,  takes  carbon 
dioxide  and  water,  supplies  the  energy  (as  light),  and  gives  us 
wood  and  oxygen.  And  the  wood  and  oxygen,  when  burned,  give 
us  back  the  original  substance,  and  the  equivalent  of  the  original 
energy  in  the  form  of  heat.  Thus,  the  two  sources  of  energy  turn 
out  to  be  the  same,  namely  the  sun's  rays. 

If,  instead  of  burning  the  starch  of  the  plant,  we  consume  it  as 
food,  it  goes  through  several  changes  instead  of  one.  But  the  final 
products  are  the  same,  namely  carbon  dioxide  and  moisture,  given 
off  through  our  lungs  and  skin,  and  heat  and  other  forms  of  energy 
such  as  are  developed  in  animals.  Thus,  whether  we  use  our 
muscles,  a  steam  engine,  or  a  water  turbine  to  do  work,  sunlight 
is  in  each  case  the  ultimate  source  of  the  energy  employed. 

Exercises.  —  1.  To  which  two  factors  in  the  interaction  of 
calcium  carbonate  and  hydrochloric  acid  (p.  381)  is  due  the  forward 
displacement  of  all  the  equilibria? 

2.  What  will  be  the  excess  of  pressure  inside  a  bottle  of  soda 
water  when  4  vols.  carbon  dioxide  are  dissolved  in  1  vol.  water? 

3.  What  volume  of  liquid  carbon  dioxide,  measured  at  0°,  will  be 
required  to  give  75  liters  of  the  gas  at  0°  and  760  mm.  pressure? 

4.  What  are  the  exact  relative  weights  of  equal  volumes  of 
carbon  dioxide,  carbon  monoxide,  air,  and  steam? 


CHAPTER  XXIX 
THE  HYDROCARBONS.     ILLUMINANTS.     FLAME 

THE  compounds  of  carbon  and  hydrogen  are  called  the  hydro- 
carbons. Hundreds  of  different  hydrocarbons,  containing  different 
proportions  of  the  two  elements  are  known.  The  natural  oil 
petroleum  is  a  mixture  of  many  substances  of  this  class. 

The  hydrocarbons  fall  into  several  distinct  series,  the  chief  one 
of  which  contains  methane  CH^  as  its  simplest  member.  On 
account  of  the  fact  that  certain  members  of  this  set  are  found  in 
paraffin,  it  is  commonly  known  as  the  paraffin  series.  For  the 
reason  that  in  this  series  the  carbon  has  all  its  four  valences  em- 
ployed, the  members  are  also  called  the  saturated  hydrocarbons. 

Paraffin  or  Saturated  Series  of  Hydrocarbons.  —  The  fol- 
lowing is  a  list  of  the  names,  formulae,  and  boiling-points  of  seven 
of  the  simplest  hydrocarbons  of  this  series,  and  of  two  of  the  higher 
members  of  the  series: 

Methane  CH4     b.-p.  -164°  Hexane  CeHn  b.-p      71° 

Ethane     C2H6      "      -  89.5°  Heptane  C7H16     "        99° 

Propane    C3H8      "      -  37°  Hexadecane         C16H34    "      287.5° 

Butane     C4H10     "      +     1°  "  "      m.-p.     18° 

Pentane    C5Hi2     "  35°  Pentatricontane  C35H72    "        74.7° 

After  the  first  four,  the  names  are  based  on  the  Greek  numerals 
representing  the  number  of  carbon  atoms  in  the  molecule.  Hep- 
tane is  followed  by  octane  C8Hi8,  nonane  C9H20,  decane  doH^, 
etc.  On  examining  the  formulae,  we  perceive  that,  in  each,  the 
number  of  hydrogen  atoms  is  equal  to  twice  the  number  of  carbon 
atoms  plus  two.  The  general  formula  is  therefore  CnH2n  +  2- 
The  series  illustrates  strikingly  the  law  of  combining  weights  (p. 
42).  We  note,  also,  that  the  first  four  are  gases  at  room  tempera- 
ture. The  members  from  pentane  to  pentadecane  Ci5H32  are 
liquids,  and  from  hexadecane  onwards  they  are  solids. 

In  these  compounds  the  carbon  is  quadrivalent,  and  each  sub- 

389 


390  COLLEGE    CHEMISTRY 

stance  is  related  to  the  preceding  one  by  containing  the  additional 
units  CH2.  The  graphic  formulae  of  the  first  three  members 
illustrate  these  two  facts: 

H  H      H  H      H     H 

I  II  III 

H-C-H  H-C-C-H  H-C-C-C-H 

I  II  III 

H  H      H  H      H     H 

These  hydrocarbons  are  extremely  indifferent  in  their  chemical 
behavior.  They  have  none  of  the  properties  of  acids,  bases,  or 
salts.  The  halogens,  notably  chlorine  and  bromine,  however,  in- 
teract with  them  (see  below).  When  burned  they  all  produce 
carbon  dioxide  and  water. 

Petroleum.  —  Petroleum  is  a  thick,  often  greenish-brown 
colored  oil.  When  borings  reach  the  oil-bearing  strata,  the  oil, 
hitherto  held  beneath  impervious  strata,  and  often  under  hydro- 
static pressure  of  water  underneath  or  around  it,  either  gushes  up 
or  is  pumped  to  the  surface.  Wells  are  in  operation  in  Caucasia, 
Gallicia,  India,  Japan,  and  in  Ontario,  Ohio,  Pennsylvania,  Cali- 
fornia and  elsewhere  in  North  America.  The  world's  production 
in  1912  was  350  million  barrels  (42  gal.  each),  of  which  nearly  220 
millions  were  produced  in  the  United  States. 

The  oil  is  a  complex  mixture,  and  is  partially  separated  by  dis- 
tillation (p.  93)  into  products  which  are  still  mixtures,  but  are 
suited  to  special  purposes.  The  components  of  lower  boiling- 
point  come  off  first  and  the  temperature  rises  steadily  as  these 
components  are  eliminated  and  those  of  higher  and  higher  boiling- 
point  enter  the  vapor.  As  certain  temperatures  are  reached  (or 
as  the  sp.  gr.  of  the  distillate  attains  certain  values)  the  condensed 
liquid  is  diverted  into  different  vessels,  so  as  to  collect  together 
the  " fractions"  of  the  same  kind.  This  is  called  fractional  dis- 
tillation. 

At  some  suitable  stage,  the  residual  oil  is  chilled,  and  a  quantity 
of  the  solid  members  of  the  series  (C22H46  to  C28H58)  crystallizes  in 
flakes  (solid  paraffin)  and  is  separated  by  filtration  in  presses. 
The  final  residue  is  used  for  lubricants  and  for  fuel.  The  fractions 
are  still  mixtures,  but  contain  mainly  compounds  lying  close  to- 
gether in  the  series.  Some  of  the  products  are  as  follows: 


THE  HYDROCARBONS 


391 


Name. 

Components. 

B.-P. 

Uses. 

Petroleum  ether 
Gasoline  .... 
Naphtha  .... 
Benzine    .... 
Kerosene     .    .    . 

Pentane,  hexane 
Hexane,  heptane 
Heptane,  octane 
Octane,  nonane 
Decane-hexadecane 

40°-  70° 
70°-  90° 
80°-120° 
120°-150° 
150°-300° 

Solvent,  gas-making 
Solvent,  fuel 
Solvent,  fuel 
Solvent 
Illuminating-oil 

Petrolatum  (vaseline),  C22H46  to  C^Kig,  is  separated  in  some 
refineries.  Solid  paraffin  is  employed  in  waterproofing  paper,  as 
an  ingredient  in  candles,  in  the  laundry,  and  to  cover  preserves. 
Kerosene,  for  oil  lamps,  is  usually  the  largest  fraction.  To  be 
used  safely,  it  should  not  give  any  inflammable  vapor  below  65° 
(150°  F.),  which  is  the  legal  flash-point  in  many  states.  Special 
treatment,  such  as  superheating  the  vapor  under  high  pressure 
(Rittman's  process),  is  used  to  increase  the  proportion  of  gasoline 
(petrol)  for  which  there  is  a  large  and  increasing  demand. 

Asphalt,  a  natural  mixture  of  solid  hydrocarbons  found  particu- 
larly in  Trinidad,  is  used  in  road-making. 

Methane  CH*.  —  Methane  is  the  chief  component  of  natural 
gas  (over  90  per  cent),  which,  like  the  oil,  is  confined  beneath 
impervious  strata  and  is  forced  out  through  borings  by  hydro- 
static pressure.  It  is  found  mainly  in  or  near  the  localities  where 
oil  is  found.  It  also  rises  to  the  surface  when  the  bottoms  of 
marshy  pools  are  disturbed  (Marsh-gas),  and  issues  from  seams  in 
coal  beds  as  fire-damp  (Ger.  Dampf,  vapor).  In  these  two  cases 
it  results  from  the  decomposition  of  vegetable  matter  in  absence 
of  air.  Its  formation  by  direct  union  of  carbon  and  hydrogen  has 
already  been  discussed  (p.  378). 

Methane  may  be  made  from  inorganic  materials  by  the  action  of 
water  upon  aluminium  carbide,  prepared  by  the  interaction  of 
aluminium  oxide  and  carbon  in  the  electric  furnace  (cf.  pp.  378, 

377): 

A14C3  +  12H2O  ->  4A1(OH)3  +  3CH4. 

In  the  laboratory  the  gas  is  commonly  obtained  by  the  distillation 
of  a  dry  mixture  of  sodium  acetate  and  sodium  hydroxide: 

NaCO2CH3  +  NaOH  ->  Na2CO3  +  CH4. 


392  COLLEGE    CHEMISTRY 

As  regards  chemical  properties,  methane,  like  other  saturated 
hydrocarbons  (p.  390),  is  very  indifferent.  When  a  mixture  of 
methane  and  chlorine  is  exposed  to  sunlight  several  changes  occur 
in  succession  (cf.  p.  162)  : 

CH4  +  C12  ->  CH3C1  +  HC1,     CH3C1  +  C12  ->  CH2C12  +  HC1, 
CH2C12  +  C12  -»  CHCU  +  HC1,    CHC13  +  C12  ->  CCU  +  HC1. 


This  kind  of  interaction  with  the  halogens  is  characteristic  of  com- 
pounds of  hydrogen  and  carbon.  It  takes  place  slowly,  and  is 
therefore  entirely  different  from  ionic  chemical  change.  It  con- 
sists, in  a  progressive  substitution  of  chlorine  for  hydrogen,  unit  by 
unit.  Chloroform  CHClj  and  carbon  tetrachloride  CCU  (p.  379) 
are  familiar  substances.  The  iodine  derivative,  iodoform  CHI3  is 
used  hi  surgical  dressing.  These  substances  are  not  salts,  and  are 
not  ionized  in  solution.  They  are  very  slowly  hydrolyzed  by 
water  —  carbon  tetrachloride,  for  example,  giving  carbonic  acid 
and  hydrochloric  acid. 

Methane  arid  the  other  saturated  hydrocarbons  are  decomposed 
by  strong  heating  (see  cracking,  below). 

Unsaturated  Hydrocarbons.  —  In  addition  to  the  saturated 
series  of  hydrocarbons,  several  other  series  are  known  in  which 
smaller  proportions  of  hydrogen  are  present.  Thus,  ethylene 
C2H4,  to  which  illuminating  gas  largely  owes  the  luminosity  of  its 
flame,  belongs  to  a  series  CnH2n,  all  the  members  of  which  contain 
two  atoms  of  hydrogen  less  than  the  corresponding  compounds  of 
the  first  series.  Again,  acetylene  C2H2  is  the  first  member  of  a  series 
CnH2n_2,  and  benzene  C6H6  begins  a  series  CnH2n_6,  of  which  toluene 
CrH8  (p.  349)  is  the  second  member.*  These  are  all  unsaturated 
because  the  full  valence  of  the  carbon  is  not  in  use,  and  these 
compounds,  therefore,  unite  more  or  less  readily  with  hydrogen, 
chlorine,  bromine,  and  concentrated  sulphuric  acid.  The  hydro- 
carbons of  all  the  series  are  mutually  soluble,  but  none  of  them 
dissolve  in  water. 

Members  of  the  ethylene  and  acetylene  series  are  found  in 

*  Isoprene  C6H8,  a  member  of  the  unsaturated  series  CnH2n-2,  when  heated 
in  presence  of  sodium  (or  some  other  contact  agent),  changes  into  caoutchouc 
(C6H8)X  or  raw  rubber.  No  method  of  preparing  synthetic  rubber  has  yet  been 
used  commercially. 


THE   HYDROCARBONS  393 

petroleum,  and  are  formed  also  to  some  extent  by  decomposition 
during  the  distillation.  As  oil  containing  them  acquires  dark- 
colored  products  by  chemical  change,  the  oils  are  always  refined 
before  being  sold.  They  are  agitated  with  concentrated  sulphuric 
acid,  which  unites  with  the  unsaturated  substances  and,  being 
insoluble  in  the  oil,  collects  in  a  layer  below  it.  The  oil  is  finally 
washed  free  from  the  acid  with  dilute  alkali  and  with  water. 

Ethylene  C3#4.  —  Ethylene  is  the  first  member  of  the  second 
series  of  hydrocarbons.  It  corresponds  to  ethane  C2H6,  but  con- 
tains in  each  molecule  two  hydrogen  units  less  than  does  this 
substance. 

Ethylene  is  made  by  heating  common  alcohol  (ethyl  alcohol) 
with  concentrated  sulphuric  acid: 


A  comparison  of  the  graphic  formulae  of  the  alcohol  and  ethylene 
shows  that  this  loss  of  water  leaves  the  carbon  partly  unsaturated: 

H      H  H      H                       H    H 

II  II                          II 

H-C-C-O-H  -C-C-     or    H-C  =  C-H 

II  II 

H      H  H      H 

The  elements  of  water  may  also  be  removed  by  allowing  the  alcohol 
to  fall  drop  by  drop  onto  heated  phosphoric  anhydride. 

Ethylene  is  formed,  along  with  acetylene  and  other  substances, 
when  any  saturated  hydrocarbon  is  heated  strongly.  Even 
methane  gives  it: 

2CH4  -»  C2H4  +  2H2. 


Ethylene  is  a  gas.  It  burns  in  the  air  with  a  flame  which,  on 
account  of  the  great  separation  of  free  carbon  which  takes  place 
temporarily  during  the  combustion  (cf.  Flame),  is  highly  luminous. 
It  will  be  seen  that,  in  the  formula,  but  three  of  the  valences  of  each 
carbon  unit  are  occupied:  the  substance  is  unsaturated.  Hence, 
when  ethylene  is  passed  through  liquid  bromine  it  is  rapidly 
absorbed,  and  the  bromine  seems  to  increase  in  volume  and  finally 
loses  all  its  color,  being  converted  into  a  transparent  liquid  having 
the  composition  C2H4Br2,  ethylene  bromide. 


394  COLLEGE    CHEMISTRY 

Acetylene.  —  This  substance,  likewise  a  gas,  is  the  first  member 
of  still  another  unsaturated  series.  Its  formula  C2H2  shows  that 
its  molecule  lacks  four  of  the  hydrogen  units  necessary  to  the  com- 
plete saturation  which  we  find  in  ethane.  Graphically  its  structure 
is  usually  represented  thus:  H  —  C  =  C  -  H.  This  gas  is 
formed  in  small  quantities  by  direct  union  of  carbon  and  hydrogen 
in  the  electric  arc  (p.  378).  This  is  because  the  reaction  is  en- 
dothermal  (p.  189).  For  the  same  reason,  it  is  also  produced  when 
ethylene  is  passed  through  a  heated  tube:  C2Hi  — >  C2H2  +  II2 
(cf.  Flame). 

When  calcium  carbide  (p.  379)  is  thrown  into  water  it  is  hy- 
drolyzed.  Violent  effervescence  occurs,  the  calcium  carbide  is 
disintegrated,  a  precipitate  of  calcium  hydroxide  is  formed,  and 
acetylene  passes  off  as  a  gas: 

CaC2  +  2H20  -»  Ca(OH)2  +  C2H2. 

This  action  is  like  that  of  water  on  calcium  phosphide  (p.  365), 
magnesium  nitride  (p.  339),  and  aluminium  carbide  (p.  391). 

Acetylene  burns  with  a  flame  which  is  still  more  luminous  than 
that  of  ethylene.  On  account  of  the  large  amount  of  heat  absorbed 
when  it  is  formed:  2C  -f  H2  -» C2H2  -  53,200  cal.,  an  equal 
amount  is  liberated  when  it  decomposes.  If  the  gas  is  compressed 
in  tanks,  it  is  therefore  apt  to  explode  from  any  shock.  It  is 
frequently  made  in  generators,  as  needed,  by  the  foregoing  action, 
and  is  used  for  lighting  on  automobiles  and  in  regions  remote  from 
a  public  supply  of  illuminating  gas.  The  acetylene  tanks,  which 
are  also  in  use,  contain  acetylene  dissolved  under  high  pressure 
in  acetone,  a  form  in  which  it  can  be  handled  safely. 

When  acetylene  C2H2  is  burned,  we  obtain  from  2  X  12  +  2  = 
26  g.  not  only  the  heat  due  to  the  combustion  of  the  carbon  (2  X 
12  X  8040  cal.,  p.  376),  and  of  the  hydrogen  (2  X  28,800  cal.),  but 
also  the  heat  due  to  the  decomposition  of  the  gas  (53,200  cal.). 
The  temperature  of  the  flame  is  therefore  extraordinarily  high. 
The  oxy acetylene  flame,  produced  by  means  of  a  suitable  burner 
(Fig.  32,  p.  58),  is  now  used,  under  the  name  of  the  acetylene 
torch,  for  cutting  metals.  The  gases  are  contained  in  portable 
tanks.  Such  a  flame  will  melt  its  way  through  a  6-inch  shaft  or  a 
steel  plate  several  feet  wide  in  less  than  a  minute,  cutting  the  object 
in  two.  Steel  buildings  have  thus  been  taken  down,  and  ships 


THE   HYDROCARBONS  395 

(like  the  Maine)  have  been  cut  up  for  removal.  Other  gases,  like 
blau  gas  and  oil  gas,  made  by  cracking  petroleum  (see  below),  are 
now  displacing  acetylene  for  this  purpose,  as  they  are  almost  as 
effective,  and  the  flame  is  more  easily  controlled. 

Cracking  of  Hydrocarbons.  —  All  hydrocarbons,  when  heated 
strongly  (air-excluded)  decompose  or  crack.  The  changes  seem 
to  be  reversible,  and  the  result  therefore  depends  upon  the  con- 
ditions. Thus,  at  atmospheric  pressure,  and  especially  when  the 
oil  is  mainly  present  as  a  liquid,  hydrogen  is  given  off  and  un- 
saturated  liquid  and  gaseous  hydrocarbons  are  produced.  Under 
such  conditions,  ethylene  is  formed  in  large  amounts.  On  the 
other  hand,  when  an  oil  free  from  gasoline  is  completely  vaporized 
(500°),  and  is  under  high  pressure,  the  hydrogen  is  forced  into 
combination  with  the  broken  molecules  and  the  saturated  con- 
stituents of  gasoline  are  produced  in  large  amount  (Rittman's 
process) . 

At  a  white  heat,  all  the  hydrocarbons  decompose  into  hydrogen 
and  free  carbon.  The  latter  is  deposited  in  a  dense  form  called 
gas-carbon,  which  is  more  or  less  crystalline  (like  graphite)  and 
used  in  making  carbon  rods  for  arc  lights  and  electric  furnaces, 
and  carbon  plates  for  batteries,  and  for  the  electrodes  employed 
in  electrolysis.  The  carbon  is  ground  up,  moistened  with  petro- 
leum residues,  subjected  to  hydraulic  pressure  and  finally  heated 
strongly  to  expel  volatile  matter. 

Carburetted  Water  Gas.  —  As  we  have  seen,  water  gas  is 
essentially  H2  +  CO  (p.  386),  and  burns  with  a  pale-blue  flame. 
To  fit  it  for  use  as  illuminating  gas,  unsaturated  hydrocarbons, 
which  burn  with  a  luminous  flame,  such  as  ethylene  Cal^  and 
acetylene  C2H2  must  be  added.  The  gas  is  sent  through  a  tower 
containing  strongly  heated  brick  on  which  a  petroleum  oil  is 
sprayed.  Mixed  with  the  vapor,  the  gas  then  passes  into  the 
"superheater"  where,  at  a  higher  temperature,  the  cracking  into 
unsaturated  hydrocarbons  occurs.  The  gas  is  then  cooled  and 
washed  to  remove  condensible  hydrocarbons,  which  would  other- 
wise obstruct  the  service  pipes.  A  typical  carburetted  water  gas 
has  the  composition:  Illuminants  17  per  cent;  heating  gases, 
methane  20  per  cent,  hydrogen  32  per  cent,  carbon  monoxide  26 


396  COLLEGE    CHEMISTRY 

per  cent;   impurities  (nitrogen  and  carbon  dioxide)  5-6  per  cent. 
A  flame  burning  5  cu.  ft.  per  hour  gives  25  candle  power. 

Blau  gas  and  Oil  gas,  such  as  Pintsch  gas,  contain  larger  pro- 
portions of  illuminants.  Thus  a  good  oil  gas  shows:  illuminants 
45  per  cent;  heating  gases,  methane  39  per  cent,  hydrogen  14.5 
per  cent;  impurities  1.5  per  cent;  candle  power  65.  Such  gases 
are  compressed  in  tanks  and  used  for  illumination  on  railway 
trains  (Coal  gas,  see  p.  410). 

FLAME 

Meaning  of  the  Term.  —  In  the  combustion  of  charcoal  there 
is  hardly  any  flame,  for  the  light  emanates  almost  entirely  from  the 
incandescent,  massive  solid.  When  two  gases  are  mixed  and  set  on 
fire,  a  sort  of  flame  passes  through  the  mixture,  but  this  can  hardly 
be  accounted  a  flame,  in  the  ordinary  sense,  either.  The  rapid 
movement  of  the  flash,  and  the  explosion  which  accompanies  it, 
are  in  a  manner  the  precise  opposite  of  the  quiet  combustion  which 
is  characteristic  of  flames. 

With  illuminating  gas  the  production  of  its  very  characteristic 
flame  is  due  to  the  chemical  union  of  a  stream  of  one  kind  of  gas 
in  an  atmosphere  of  another.  The  flame  is  made  up 
of  the  heated  matter  where  the  two  gases  meet.  In 
the  case  of  a  burning  candle  (Fig.  102),  one  of  the 
bodies  appears  to  be  a  solid,  but  a  closer  scrutiny 
of  the  phenomenon  shows  that  the  solid  does  not 
burn  directly.  A  combustible  gas  is  manufactured 
continuously  by  the  heat  of  the  combustion  and 
rises  from  the  wick.  The  introduction  of  a  narrow 
tube  into  the  interior  of  the  flame  enables  us  to 
Fio.  102.  draw  off  a  stream  of  this  gas  and  to  ignite  it  at  a 
remote  point.  Thus,  a  flame  is  a  phenomenon  produced  at  the 
surface  where  two  gases  meet  and  undergo  combination  with  the 
evolution  of  heat  and,  more  or  less,  light. 

In  the  chemical  point  of  view,  it  is  a  matter  of  indifference 
whether  the  gas  outside  the  flame  contains  oxygen,  and  the  gas 
inside  consists  of  substances  ordinarily  known  as  combustibles,  or 
whether  this  order  is  reversed.  In  an  atmosphere  of  ordinary 
illuminating  gas,  the  flame  must  be  fed  with  air.  This  condition 


FLAME  397 

is  easily  realized  (Fig.  103).  The  lamp  chimney  is  closed  at  the 
top  until  it  has  become  filled  with  illuminating  gas.  This  can  be 
ignited  as  it  issues  from  the  bottom  of  the  wide,  straight  tube. 
When  the  hole  in  the  cover  of  the  lamp  chimney 
is  then  opened,  the  upward  draft  causes  the  flame 
of  the  burning  gas  to  recede  up  the  tube,  and  there 
results  a  flame  fed  by  air  and  burning  in  coal  gas. 

Luminous  Flames.  —  The  flame  of  hydrogen, 
under  ordinary  circumstances,  is  almost  invisible, 
nearly  all  the  energy  of  the  combustion  being  de- 
voted to  the  production  of  heat.  A  part  of  this, 
however,  may  be  transformed  into  light  by  the 
suspension  of  a  suitable  solid  body,  such  as  a 
platinum  wire,  in  the  flame.  The  holding  of  a 
piece  of  quicklime  in  an  oxy hydrogen  flame  (cf.  p. 
58)  is  a  practical  illustration  of  this  method  of 
securing  luminosity.  In  general,  luminosity  may 
be  produced  by  the  presence  of  some  solid  which  is  heated  to 
incandescence. 

In  the  Welsbach  lamp  the  flame  itself  is  nonluminous  and,  but 
for  the  mantle,  would  be  identical  with  the  ordinary  Bunsen  flame. 
The  mantle  which  hangs  in  the  flame,  however,  by  its  incandes- 
cence, furnishes  the  light.  This  mantle  is  composed  of  a  mixture 
of  99  per  cent  thorium  dioxide  Th02  and  one  per  cent  cerium 
dioxide  Ce02.  These  oxides  act  as  a  contact  agent,  hastening 
the  combustion  and  liberation  of  heat  close  to  their  surface,  and  so 
acquire  a  temperature  higher  than  the  average  for  the  rest  of  the 
flame.  The  Welsbach  lamp  gives  four  tunes  as  much  light  as  does 
the  same  gas,  issuing  at  the  same  rate,  from  an  ordinary  burner. 

In  cases  of  brilliant  combustion,  as  of  magnesium  ribbon  or 
phosphorus,  a  solid  body  is  formed  whose  incandescence  accounts 
for  the  light.  The  flame  of  ordinary  illuminating  gas  does  not  at 
first  sight  appear  to  give  evidence  of  the  presence  of  any  solid  body. 
But  if  a  cold  evaporating  dish  is  held  in  the  flame  for  a  moment,  a 
thick  deposit  of  finely  divided  carbon  (soot)  is  formed,  and  we  at 
once  realize  that  the  light  is  due  to  the  glow  of  these  particles  in  a 
mass  of  intensely  hot  gas.  Carbon  is,  indeed,  an  extremely  com- 
bustible substance,  and  is  eventually  entirely  consumed.  But  a 


398  COLLEGE    CHEMISTRY 

fresh  supply  is  being  generated  continuously  in  the  interior  of  the 
flame,  while  the  oxygen  with  which  it  is  to  unite  is  outside  the 
flame  altogether.  Thus  the  carbon  particles  persist  until,  drifting 
with  the  spreading  gas,  they  reach  the  periphery  of  the  flame. 

On  a  large  scale,  oil  residues  are  burned  so  that  the  flame  strikes 
a  revolving,  iron  vessel  cooled  with  water.  The  soot  or  lampblack 
is  continuously  scraped  off  as  the  vessel  turns.  Lampblack  is 
used  in  making  printer's  ink,  India  ink,  and  black  varnish. 

The   Bunsen   Flame.  —  In   the   burner  devised   by   Robert 
Bunsen,  a  jet  of  ordinary  illuminating  gas  is  projected  from  a  ntir- 
.  row  opening  into  a  wider  tube  (Fig.  104).     In  this 

A  tube  it  becomes  mixed  with  air,  entering  through 

openings  whose  dimensions  can  be  altered  by  means 
of  a  perforated  ring.  When  the  supply  of  air  is 
sufficient,  the  flame  becomes  non-luminous.  With 
a  somewhat  different  construction,  and  the  use  of 
a  bellows  to  force  a  larger  proportion  of  air  into 
the  gas,  a  still  hotter  flame  can  be  produced.  The 
instrument  in  this  case  is  known  as  a  blast  lamp. 

The  high  temperature  of  the  blast  lamp  flame 
presents  an  interesting  problem.  The  same  amounts 
of  gas  and  air  burn  to  give  the  same  amounts  of 
the  same  products,  whether  the  air  blast  is  on  or 
off.  The  same  amount  of  heat  is  produced,  and 
the  same  quantities  of  the  same  substances  are 
heated.  The  average  temperature  throughout  the 
flame  should  therefore  be  the  same.  In  point  of  fact,  it  is  the 
same,  but  the  stream  of  hot  gas  is  moving  more  rapidly  when  the 
blast  is  going.  The  temperature  of  a  body  immersed  in  the  flame 
depends,  on  the  one  hand  upon  the  rate  at  which  heat  reaches  it, 
and  upon  the  other  on  the  rate  at  which  it  loses  heat  by  radiation. 
The  heat  is  partly  carried  by  the  moving,  heated  gases  (convection), 
and  partly  transmitted  by  conduction  through  the  stationary  layer 
(p.  331)  on  the  surface  of  the  body.  Now,  the  latter  is  the 
slower  process.  Hence  a  rapid  stream  of  gas,  which  leaves  a 
thinner  stationary  layer,  will  diminish  the  distance  the  heat  has  to 
travel  by  conduction  and  so  convey  heat  to  the  body  faster  than 
could  a  slow  stream  of  the  same  temperature.  Thus,  with  a  blast 


FLAME 


399 


flame,  the  loss  by  radiation  is  the  same  at  the  same  temperature, 
but  heat  reaches  the  body  faster  and  so  the  temperature  of  the 
body  more  nearly  approaches  that  of  the  flame  itself. 

Structure  of  the  Illuminating  and  the  Bunsen  Flame.  - 

When  an  exceedingly  small  luminous  flame  is  examined,  the  various 
parts  of  which  it  consists  may  easily  be  made  out.     In  the  ulterior 
there  is  a  dark  cone  which  is  composed  of  illuminating  gas  and  air, 
and  in  it  no  combustion  is  taking  place.     A  match4iead  may  be 
held  here  for  some  time  without  being  set  on  fire.     This  is  there- 
fore not  properly  a  part  of  the  flame.     Outside  this  is  a  vivid  blue 
layer  (C,  Fig.  105)  which  is  best  seen  in  the  lower  part 
of  the  flame,  but  extends  beneath  the  luminous  sheath, 
and  covers  the  dark  inner  cone  completely.     Outside 
the  blue  flame,  and  covering  the  greater  part  of  it,  is 
the  cone-shaped  luminous  portion  (B).     Over  all  is  a 
faint  mantle  of  non-luminous  flame  (A),  which  be- 
comes visible  only  when  the  light  from  the  luminous 
part  is  purposely  obstructed.     In  the  luminous  gas- 
flame,  therefore,  there  are  four  regions,  if  we  count 
the  inner  cone  of  gas.     The  difference  between  this 
and  the  non-luminous  Bunsen  flame  is  that  in  the 
latter  the  luminous  region  is  omitted,  and  the  inner, 
dark  cone,  the  blue  sheath,  and  the  outer  mantle, 
are  the  only  parts  which  can  be  distinguished. 

The    Causes   of  Luminosity  and  Non-Lumi-      ^ 

nosity.  —  The  study  of  the  chemical  changes  taking  ^  1Q5 
place  in  the  Bunsen  flame,  particularly  with  the 
object  of  explaining  (1)  the  luminosity  of  the  flame  of  the  pure 
gas  and  (2)  the  non-luminosity  of  that  produced  by  the  same  gas 
when  it  is  mixed  with  air,  has  been  the  subject  of  many  elaborate 
investigations.  The  questions  are:  ;  (1)  Why  *  ^™,h^™ 
in  the  former  case,  and  (2)  why  is  it  not  liberated  in  the  latter? 
Let  us  consider  these  questions  in  order. 

1  The  investigations  of  Lewes  (1892)  and  others  show  conclu- 
sively that  the  free  carbon  in  the  luminous  zone  of  the  ordinary 
flame  is  accompanied  by  free  hydrogen,  and  that  both  are  formed 
by  dissociation  of  the  ethylene.  This  substance,  when  heated, 


COLLEGE    CHEMISTRY 


>  H2  +  C2H2  ->  2C  +  H2. 

The  carbon  glows,  until,  as  it  drifts  outwards,  it  encounters  the 
oxygen  of  the  air  and  is  burned.  The  first  oxygen  encountered 
combines  more  readily  with  the  hydrogen,  since  it  is  a  gas,  than 
with  the  carbon,  which  is  now  in  solid  particles  and  therefore  burns 
less  readily  That  carbon  glows  when  heated  in  the  absence  of 
oxygen,  without  being  consumed,  is  a  fact  familiar  in  the  behavior 
E  the  incandescent  electric  lamp,  the  filament  of  which  is  often 
made  of  carbon. 

The  conception  that  when  hydrocarbons  burn,  they  first  undergo 
dissociation,  and  then  union  with  oxygen,  is  in  harmony  with  what 
we  have  observed  also  in  the  case  of  the  combustion  of  hydrogen 
ulphide,  where  the  presence  of  free  sulphur  and  free  hydrogen  in 
J  interior  of  the  flame  was  demonstrated  (p.  268). 
2.   The  influence  of  the  air  admitted  to  the  Bunsen  burner  in 
interfering  with  this  dissociation  in  such  a  way  as  to  destroy  all 
luminosity,  is  the  most  difficult  point  to  explain      The 
effect  is  frequently  attributed  to  the  oxygen  which  the 
air  contains.     This  view,  however,  is  seriously  weak- 
ened by  a  consideration  of  the  undoubted  fact  that 
oxygen  is  not  required.     Carbon  dioxide  and  steam 
are  equally  efficient  when  introduced  instead  of  air 
[Fig.  106,  gas  enters  at  a  and  CO2  at  6).    Even  nitro- 
gen, which  cannot  possibly  be  suspected  of  furnishing 
any  oxygen,  likewise  destroys  the  luminosity.     Lewes 
has  shown  that  0.5  volumes  of  oxygen  in  1  volume  of 
coal  gas  destroy  the  luminosity.     But  2.30  volumes 
of  nitrogen  or  2.27  volumes  of  air  accomplish  the  same 
result.     Thus  the  efficiency  of  air  is  not  much  greater 
than  that  of  nitrogen,  in  spite  of  the  fact  that  one- 
fifth  of  the  former  is  oxygen. 
It  is  evident  that  the  effect  is  due,  in  part  at  least,  to  the  dilution 
with  a  cold  gas.     This  is  confirmed  by  the  observation  that  a  cold 
platinum  dish  held  in  a  small  luminous  flame  is  similarly  destruc- 
tive of  the  luminosity.     If  the  tube  of  the  Bunsen  burner  is  heated 
so  that  the  mixed  gases  are  considerably  raised  in  temperature 


FLAME  401 

before  reaching  the  non-luminous  flame,  the  latter  becomes  lumi- 
nous. It  is  probable,  therefore,  that  the  cold  gas  lowers  the  tem- 
perature of  the  inner  flame,  and  at  the  same  time  the  dilution 
diminishes  the  speed  with  which  the  free  carbon  is  formed  (Lewes). 
Even  if  the  temperature  is  not  reduced  below  that  at  which  dis- 
sociation of  the  ethylene  can  occur,  yet  the  dilution  and  cooling, 
together,  prevent  that  sharp  dissociation  at  this  particular  point 
which  is  necessary  for  the  production  of  the  great  excess  of  free 
carbon  needed  to  furnish  the  light. 

Before  these  investigations  were  made,  a  different  answer  was 
given  to  the  question  why  the  flame  of  pure  illuminating  gas  con- 
tains free  carbon  and  is  luminous.  It  was  said  that  hydrogen  was 
more  easily  burned  than  carbon,  and  therefore  the  latter  was  left 
free,  to  be  burned  later.  It  is  true  that  gaseous 
hydrogen  burns  more  easily  than  solid  carbon,  e.g., 
charcoal.  But  in  ethylene,  both  elements  are  equally 
gaseous  and  the  explanation  is  faulty.  Smithells 
(1892)  demonstrated  the  falsity  of  this  explanation 
by  devising  a  cone-separator  (Fig.  107).  The  air 
and  ethylene  or  other  gas  are  admitted  separately, 
and  the  inner  cone  of  the  non-luminous  flame  rests 
on  the  inner,  narrow  tube,  while  the  outer  cone  is 
at  the  top.  By  means  of  a  side  tube  (not  shown)  he 
withdrew  the  inter-conal  gas  and  found  that,  while 
all  of  the  carbon  was  burned  by  the  inner  cone  as  far 
as  carbon  monoxide  CO,  most  of  the  hydrogen  was 
still  entirely  uncombined.  The  change  in  the  inner  FIQ  1Q7 
cone  of  the  Bunsen  flame  consists,  therefore,  mainly 
in  the  burning  of  all  the  hydrocarbons  to  carbon  monoxide,  with 
liberation  of  the  hydrogen.  In  the  outer  cone,  it  is  practically  a 
burning  of  water  gas  that  is  taking  place. 

Exercises.  —  1.   Write  a  graphic  formula  for  hexane. 

2.  Write  an  equation  for  the  formation  of  aluminium  carbide 

(p.  391). 

3.  Make  a  section  showing  the  shape  of  the  flame  produced  by 
burning  hydrogen  gas  when  the  latter  issues  from  a  circular 
opening. 


CHAPTER  XXX 
THE   CARBOHYDRATES  AND  RELATED   SUBSTANCES 

A  PLANT  takes  carbon  dioxide  from  the  air  and  water  from  the 
ground  and,  using  the -energy  of  sunlight,  converts  them  into  a 
growing  framework  of  cellulose  (CeHioOs),,  and,  as  we  have  seen 
(p.  387),  into  starch  (CeHioOs)^  which  it  stores  in  the  cells. 
The  cellulose  of  certain  plants  furnishes  us  with  cotton,  linen, 
jute,  and  paper.  The  starch  of  wheat,  oats,  maize  (corn),  and 
potatoes  is  one  of  the  chief  foodstuffs  they  contain.  The  plant, 
when  dead  and  buried,  changes  into  coal.  The  fresh  wood,  when 
distilled,  supplies  wood  spirit  (methyl  alcohol)  and  other  useful 
substances,  and  the  residue  is  the  valuable  charcoal.  Further- 
more, from  starch  we  can  readily  make  sugar,  alcohol,  and  other 
familiar  materials.  Cellulose,  starch,  and  the  sugars  (e.g.,  cane- 
sugar  C^HjaOii)  contain  oxygen  and  hydrogen  in  the  same  pro- 
portions in  which  they  are  present  in  water,  2H  :  1O.  They  might 
be  considered  hydrates  of  carbon,  and  so  they  are  called  the 
carbohydrates.  The  foregoing  brief  summary  shows  that  the 
carbohydrates  introduce  us  to  a  much  greater  variety  of  inter- 
esting organic  compounds  than  does  petroleum. 

Cellulose  (C6#1006)*  and  Paper.  — The  wall  of  each  cell, 
and  therefore  the  whole  framework  of  a  plant,  is  made  of  cellulose. 
Linen  and  cotton  are  pure  cellulose.  The  walls  of  the  cells  are 
usually  more  or  less  thickened  by  a  substance  called  lignin,  which 
has  much  the  same  composition,  but  different  chemical  behavior. 
The  best  paper  is  made  of  cotton  or  linen  (rag-paper).  Cheaper 
kinds  are  prepared  by  cutting  wood,  such  as  spruce  or  pine,  into 
chips  and  treating  (" cooking")  them  with  a  solution  of  calcium 
bisulphite  Ca(HSO3)2.  This  process  decomposes  the  lignin,  and 
converts  it  into  soluble  materials.  The  sulphite  liquor  is  then 
run  off,  and  the  pulpy  material  is  washed,  beaten  with  water  to 
reduce  it  to  minute  shreds,  and  bleached  with  very  dilute  chlorine- 

402 


THE    CARBOHYDRATES   AND   RELATED   SUBSTANCES      403 

water.  The  pure  cellulose,  now  paper-pulp,  is  suspended  in 
water,  spread  on  screens,  pressed,  and  dried.  During  the  process, 
other  substances  are  added.  Thus,  size  (gelatine  or  rosin  and 
alum,  see  Sizing)  prevents  the  ink  from  running;  pulverized 
calcium  sulphate  (gypsum),  clay,  and  other  white  solids  ("load- 
ing") give  body  to  the  paper  and  permit  the  production  of  a 
smooth  surface  by  rolling  ("calendering").  Dyestuffs  can  be 
added  to  give  special  tints.  Filter  paper  is  pure  cellulose. 

Starch  (C6H10  Os)y.  — Starch  consists  of  little  colorless  granules 
of  various  rounded  shapes  (Fig.  108),  which  are  easily  seen  under 
the  microscope.  These  granules  are  massed  in  large  quantities 
in  the  ears  of  wheat  and  oats,  in  the 
tubers  of  potatoes,  in  the  grams  of 
maize  (corn),  and  in  peas  and  beans. 
Even  in  the  leaves  they  can  be  seen. 
Starch  is  recognized  by  the  iodine  test 
(p.  3),  turning  deep  blue  with  a  trace 
of  free  iodine.  ^^ 

The  treatment  of  wheat  flour,  which 

is  three-fourths  starch,  by  washing  out  the  starch  through  a  porous 
cloth  with  water,  has  already  been  described  (p.  3).  It  is  made 
from  maize  in  America  and  from  potatoes  in  Europe,  by  washing 
the  flour  on  sieves. 

Starch  is  not  soluble  in  water.  If  boiled  with  water,  however, 
the  granules  swell  and  break  and  the  starch  is  diffused  through 
the  water,  giving  a  clear  liquid.  If  too  much  water  is  not  used, 
the  liquid  when  cold  sets  as  a  jelly.  While  the  liquid  is  hot,  much 
of  the  starch  will  pass  through  a  filter  along  with  the  water. 
Such  a  liquid  is  called  a  colloidal  suspension.  Suspensions  like 
this  are  constantly  met  with  in  using  complex  organic  compounds 
like  jellies,  glues,  soaps,  and  dyes.  Even  insoluble  inorganic 
materials,  like  gold,  give  such  suspensions  (see  p.  416). 

The  colloidal  suspension  of  starch  turns  blue  when  a  solution 
containing  free  iodine  is  added  to  it.  It  is  used  in  the  laundry  for 
stiffening  white  goods.  Glucose  is  manufactured  from  it. 

Glucose  Ce#12O6,  a  Sugar,  from  Starch.  —  When  starch  is 
boiled  with  water,  to  which  a  few  drops  of  an  acid  (contact  agent), 


404  COLLEGE    CHEMISTRY 

such  as  hydrochloric  acid,  have  been  added,  the  liquid,  after 
neutralization  of  the  acid,  is  found  to  be  sweet.  One  of  the  sugars, 
glucose  CgH^Oe,  can  be  obtained  in  crystals  after  evaporation. 
The  crystals  form  "brewers'  glucose"  and  the  syrup  produced  by 
concentration  is  corn  syrup  (if  maize  is  the  source  of  the  starch). 
The  latter,  although  less  sweet  than  ordinary  sugar,  is  much  less 
expensive  and  is  used  in  making  preserves  and  cheap  candy. 

The  molecular  weight  of  starch  is  unknown,  but  undoubtedly 
large.  The  formula  (CeHuAOi,  shows  the  composition.  The 
water,  in  presence  of  a  little  acid,  decomposes  the  molecules  and 
combines  with  the  material.  First,  dextrin  (used  as  paste  or 
mucilage)  is  formed  and  this  breaks  up  into  glucose.  The  action 
is  an  hydrolysis: 

B),  +  2/H20 


Glucose  is  known  also  as  dextrose  and  as  grape  sugar.  The 
crystalline  granules  in  raisins  (dried  grapes)  are  composed  mainly 
of  it.  When  pure,  it  is  almost  colorless.  It  reduces  cupric 
hydroxide,  in  Fehling's  solution  (q.v.),  to  cuprous  oxide. 

Corn  syrup  contains  30-40  per  cent  of  unchanged  dextrin,  40-50 
per  cent  of  glucose,  and  the  rest  water. 

The  Sugars.  —  The  common  sugars  may  be  divided  into  the 
monosaccharides,  usually  with  the  formula  CeH^Oe,  and  the 
disaccharides,  usually  C^H^On.  Of  these,  the  following  will  be 
referred  to  in  what  follows: 


Monosaccharides:  Glucose  (grape  sugar  or  dextrose) 

Fructose  (fruit  sugar  or  levulose)  CeH^Oe. 

Disaccharides:         Sucrose  (cane-sugar,  beet-sugar)  C^H^On. 
Maltose  (action  of  malt  on  starch)  C^H^On. 
Lactose  (milk  sugar,  in  animals  only)  C^H^On. 

Sucrose,  or  Cane  Sugar.  —  Plants,  such  as  the  sugar-cane 
and  beet,  besides  forming  cellulose  and  starch,  produce  excep- 
tional amounts  of  sucrose,  or  table  sugar.  The  sap  of  the  sugar 
maple  contains  much  of  it. 

Cane  sugar  is  extracted  by  crushing  the  stalks  between  rollers, 
and  evaporating  the  expressed  liquid  (18  per  cent  sugar)  in  closed 
pans.  A  partial  vacuum  is  maintained  so  that  the  solution  may 


THE    CARBOHYDRATES   AND   RELATED    SUBSTANCES      405 

boil  at  a  low  temperature  (65°  to  begin  with)  and  none  of  the 
sugar  be  decomposed.  When  the  syrup  cools,  the  sucrose  ap- 
pears in  brown-colored  crystals.  The  mother-liquor  is  called 
molasses.  In  the  refinery,  the  sugar  is  redissolved,  the  solution  is 
poured  through  a  column  of  charcoal,  which  takes  out  the  coloring 
matter,  and  the  liquid  is  once  more  allowed  to  deposit  crystals. 
Pure  cane-sugar  has  a  faint  yellow  tint,  and  a  small  amount  of 
ultramarine  (q.v.)  is  added  to  give  that  whiteness  which  is  pop- 
ularly connected  with  purity  in  sugar. 

Sugar  beets  (16  per  cent  or  more  sugar)  are  sliced  and  steeped 
in  water.  The  extract  contains  a  gummy  material  in  colloidal 
suspension.  This  is  coagulated  and  precipitated  by  adding  slaked 
lime  (calcium  hydroxide)  Ca(OH)2  suspended  in  water,  and  boil- 
ing. After  separation  of  the  clear  liquid,  carbon  dioxide  is  passed 
through  it  to  precipitate  the  excess  of  lime  as  carbonate  (CaC03). 
The  solution  is  then  decolorized  with  charcoal  and  evaporated  to 
crystallization. 

As  regards  properties,  sucrose  crystallizes  in  four-sided  prisms 
(rock-candy,  Fig.  48,  p.  83)  and  melts  at  160°.  When  heated  to 
200-210°  it  partially  decomposes,  leaving  a  soluble,  brown,  mixed 
material,  caramel,  used  in  coloring  whisky  and  soups.  Sucrose 
does  not  reduce  Fehling's  solution. 

When  boiled  with  water  containing  a  trace  of  acid  (contact 
agent),  sucrose  is  hydrolyzed,  giving  a  mixture  of  the  two  mono- 
saccharides,  glucose  and  fructose: 

C^H^Oii  -f-  H^O  —  >  CeH^Oe  ~h 


The  mixture  is  called  invert  sugar,  and  is  found  in  many  sweet 
fruits  and  in  honey.  Each  sugar  interferes  with  the  crystalliza- 
tion of  the  other,  by  lowering  the  freezing  point  (p.  134),  and  so 
invert  sugar  is  added  in  making  fondant  candy  and  candy  that 
is  to  be  pulled,  both  of  which  must  remain  soft  for  some  time. 
Icing  for  cakes  has  to  some  extent  this  property,  and  is  made  by 
adding  acid  substances  to  sugar,  such  as  vinegar,  lemon  juice,  or 
cream  of  tartar. 

Enzymes.  —  Yeast,  consisting  of  microscopic  cells,  belongs  to 
a  low  order  of  plants.  Its  use  lies  in  the  fact  that,  while  multi- 
plying, it  secretes  within  each  cell  two  very  active,  soluble  sub- 


406 


COLLEGE    CHEMISTRY 


stances.  These  are  zymase  and  sucrase  (invertase)  which  belong 
to  a  class  of  organic  substances  called  enzymes.  Sucrase  means 
an  enzyme  that  splits  sugar.  Enzymes  produce  remarkable 
chemical  changes  by  their  mere  presence  (contact  actions). 

Alcoholic  Fermentation.  —  When  some  yeast,  which  is  a 
mass  of  the  living  plants,  is  added  to  a  solution  of  glucose  at 
about  30°,  the  small  amount  of  zymase  gradually  decomposes 
the  sugar.  Bubbles  of  carbon  dioxide  soon  begin 
to  rise,  and  may  be  tested  (p.  384)  with  limewater 
(Fig.  109).  At  the  same  time,  alcohol  (ethyl  alco- 
|[  hoi  C2H5OH)  accumulates  in  the  liquid: 

->  2CO2 1  +  2C2H5OH. 


FIG.  109. 


Yeast  will  ferment  fructose  CeH^Oe  with  the  same 
result,  but  more  slowly,  so  that,  when  placed  in 
invert  sugar,  it  decomposes  the  glucose  first  and 
the  fructose  afterwards. 

Zymase  does  not  act  upon  sucrose  (table  sugar),  but  sucrase 
hydrolyzes  the  sucrose  in  the  same  way  as  does  a  dilute  acid,  giv- 
ing invert  sugar.  The  latter  is  then  decomposed  by  the  zymase, 
and  so  cane-sugar  in  solution  is  fermented  by  yeast  into  alcohol 
and  carbon  dioxide,  just  as  is  glucose,  only  more  slowly. 

In  making  wine  the  glucose  contained  in  the  grape  juice  is  fer- 
mented by  a  species  of  yeast  found  on  the  skins  of  the  grapes. 

Commercial  Alcohol  is  not  made  from  sugar,  but  from  the  starch 
of  potatoes  or  maize.  When  barley  is  allowed  to  sprout,  an 
enzyme,  amylase  (meaning  starch-splitting  enzyme)  or  diastase, 
is  formed  in  the  ears.  The  whole  material  is  dried,  and  is  then 
called  malt.  When  this  is  mixed  with  starch  and  water,  the 
amylase  hydrolyzes  the  starch  to  maltose  C^H^On  (p.  404). 
The  latter  is  then  further  hydrolyzed  by  yeast  to  form  glucose 
CeHtfOe,  and  this  is  decomposed  by  the  zymase  into  alcohol  and 
carbon  dioxide. 

Whisky  (about  50  per  cent  alcohol)  is  made  by  treating  the 
starch  of  rye,  maize  or  barley  in  the  same  way,  with  subsequent 
distillation  (see  below)  to  separate  the  alcohol  (whisky).  Beer  is 
made  similarly  from  various  kinds  of  grain,  especially  barley,  but 
the  fermented  liquor  is  not  distilled. 


THE   CARBOHYDRATES   AND    RELATED   SUBSTANCES      407 

Ethyl  Alcohol  CQH5OH.  —  Common  alcohol  is  related  to 
ethane  C2H6,  having  an  hydroxyl  group  in  place  of  one  unit  of 
hydrogen.  Hence  its  name,  ethyl  alcohol. 

Ethyl  alcohol  boils  at  78.3°  and  so,  when  the  fermented  liquor 
is  distilled,  it  is  almost  pure  alcohol  that  comes  off.  Commercial 
alcohol  contains  95  per  cent  by  volume  (in  Great  Britain,  90  per 
cent).  Absolute  alcohol  is  made  by  adding  quicklime,  which 
combines  with  the  water,  and  redistilling  the  liquid. 

Alcohol  mixes  with  water  in  all  proportions.  In  dilute,  aqueous 
solution  it  is  not  ionized,  and  does  not  interact  with  acids,  bases, 
or  salts.  It  is,  however,  easily  oxidized  to  acetic  acid.  When 
water  is  absent,  it  interacts  with  acids  slowly  (see  p.  413). 

Alcohol  is  used  as  a  solvent  for  the  resins  employed  in  making 
varnishes  for  wood  and  lacquers  for  metal. 

On  account  of  the  high  duty  on  95  per  cent  alcohol  ($2.11  per 
gallon  in  the  U.  S.  and  24/6  in  Gt.  Britain),  denatured  alcohol 
(methylated  spirit),  which  is  free  of  duty,  is  employed  for  indus- 
trial purposes.  The  alcohol  (cost  about  22  cents  per  gal.  in  the 
U.  S.  and  1/6  in  Britain)  is  mixed  with  offensive  or  poisonous 
materials,  which  prevent  its  consumption  as  a  beverage,  without 
interfering  with  other  uses.  Wood  spirit  and  gasoline  are  often 
employed. 

Acetic  Acid  HCO2CH3.  —  This  is  the  sour  substance  in  vinegar, 
and  has  many  industrial  applications.  Vinegar  is  made  by  oxi- 
dizing alcohol  with  atmospheric  oxygen,  using  an  enzyme  se- 
creted by  bacterium  aceti  (mother  of  vinegar)  as  contact  agent. 
Dilute  alcohol  from  any  source,  such  as  fermented  apple  juice 
(hard  cider),  is  allowed  to  trickle  over  shavings  in  a  barrel.  Holes 
admit  air,  and  the  shavings  are  inoculated  in  advance  by  wetting 
with  vinegar: 

HOC2H5  +  02  ->  HC02CH3  +  H2O. 

The  issuing  liquid  contains  5-15  per  cent  of  acetic  acid,  which  can 
be  purified  by  fractional  distillation  to  separate  the  water.  It 
boils  at  118°  and  freezes  at  16.7°.  Although  four  atoms  of  hydro- 
gen are  contained  in  its  molecule,  but  one  of  these  is  replaceable 
by  metals.  This  fact  is  recognized  in  the  reaction  formula  (p. 
95)  of  the  acid,  HC2H302,  or  HCO2CH3.  It  is  a  weak,  monobasic 
acid:  HC02CH3 <=*  H+  + 


408  COLLEGE    CHEMISTRY 

Destructive  Distillation  of  Wood.  Charcoal.  —  Dry  wood 
is  distilled  in  iron  retorts,  and  the  vapors  coming  off  are  led  through 
a  condenser  to  separate  the  liquids  from  the  gases.  The  cellulose, 
lignin,  and  resinous  material  are  decomposed,  and  only  charcoal 
remains.  The  gases,  consisting  mainly  of  hydrogen,  methane 
CHi,  ethane  C2H6,  ethylene  C2H4,  and  carbon  monoxide  CO,  are 
employed,  on  account  of  their  combustibility,  as  fuel  in  the  dis- 
tillation itself.  The  fluids  form  a  complex  mixture  containing 
large  quantities  of  water,  methyl  alcohol  CH3OH  (wood  spirit), 
acetic  acid,  acetone  (CH3)2CO,  and  tar.  The  liquids  can  be 
separated.  The  methyl  alcohol  (wood  spirit)  is  used  in  varnish 
making.  The  acetone  has  several  uses  (e.g.,  p.  394). 

Wood  charcoal  exhibits  the  cellular  structure  of  the  material 
from  which  it  was  made,  and  is  therefore  highly  porous  and  has 
an  enormous  internal  surface.  When  the  charcoal  is  burned,  the 
mineral  constituents  of  the  wood  appear  in  the  ash.  This  is 
composed  of  the  carbonates  of  the  metallic  elements  present. 
For  certain  purposes,  charcoals,  made  in  the  same  fashion  as 
the  above  from  bones  and  from  blood,  find  wide  application. 
The  former,  called  bone  black,  contains  much  calcium  phosphate 
(p.  362).  In  the  old  method  of  making  charcoal,  which  is  still 
practised,  the  wood  was  piled  up,  covered  with  turf,  and  set  on 
fire.  All  the  valuable  volatile  products  were  lost,  as  well  as  part 
of  the  charcoal  itself. 

Properties  of  Charcoal.  —  Charcoal  exhibits  certain  proper- 
ties which  are  not  shared  by  other  forms  of  carbon.  For  example, 
it  can  take  up  large  quantities  of  many  gases.  Boxwood  charcoal 
will  in  this  way  absorb  ninety  times  its  own  volume  of  ammonia, 
fifty-five  volumes  of  hydrogen  sulphide  or  nine  volumes  of  oxygen. 
Freshly  made  dogwood  charcoal  (used  in  making  the  best  gun- 
powder), when  pulverized  immediately  after  its  preparation, 
often  catches  fire  spontaneously  on  account  of  the  heat  liberated 
by  the  condensation  of  oxygen.  It  is  therefore  set  aside  for  two 
weeks,  to  permit  the  slow  absorption  of  moisture  and  air.  The 
absorbed  gases  may  be  removed  unchanged  by  heating  the  char- 
coal in  a  vacuum.  The  phenomenon,  described  as  adsorption,  is 
caused  by  the  adhesion  of  the  gases  to  the  very  extensive  surface 
(due  to  porosity)  which  the  charcoal  possesses.  Glass  and  all  other 


THE    CARBOHYDRATES  AND    RELATED   SUBSTANCES      409 

solids  show  the  same  property,  though  in  a  smaller  degree  (p.  88). 
Solid  and  liquid  bodies  are  also  in  many  cases  taken  up  by  charcoal 
in  a  similar  fashion.  Thus,  organic  dyes,  such  as  indigo,  litmus, 
and  cochineal,  and  natural  coloring  matters  (see  sugar  refining, 
p.  405),  which  are  all  more  or  less  colloidal  in  nature,  are  removed 
when  the  liquid  is  shaken  with,  or  poured  through  pulverized 
charcoal.  The  organic  materials  dissolved  in  drinking  water 
also  undergo  adsorption  in  charcoal,  but  the  charcoal  soon  be- 
comes inactive.  Charcoal  is  likewise  used  in  reducing  ores,  and 
as  a  smokeless  fuel. 

Coal.  —  When  vegetable  matter  decomposes,  without  heating, 
and  while  covered  with  sand  or  clay  so  that  air  is  excluded,  water 
and  hydrocarbons  are  liberated,  and  the  products  are  peat,  bitumi- 
nous coal,  or  anthracite. 

We  are  concerned  mainly  with  the  products  obtained  by  dis- 
tilling coal,  to  get  coal  gas  and  coke,  and  with  its  use  as  fuel.  To 
determine  its  suitability  for  various  purposes,  the  coal  is  analyzed, 
and  its  heating  power  is  measured. 

In  coal  analysis,  the  air-dried  material  is  used.  The  water  is 
determined  by  heating  1  g.  at  105°  for  1  hour.  Much  water  lowers 
the  fuel  value,  because  heat  is  wasted  in  vaporizing  it,  and  in  de- 
composing it  (cf.  p.  386).  After  re  weighing  the  sample,  the  coal  is 
heated  with  the  Bunsen  flame  in  a  covered  crucible  to  drive  off 
the  volatile  matter.  After  weighing  again,  air  is  admitted,  and 
strong  heat  is  applied  to  burn  up  the  fixed  carbon  (coke).  The 
residue  is  the  ash.  In  the  following  table  the  proportions  are  com- 
pared with  seasoned  wood  on  the  one  hand,  and  with  charcoal  and 
coke  on  the  other. 

The  calorific  power  of  a  coal  determines  largely  its  value  for 
heating.  A  sample  (about  1  g.)  is  burned  in  a  bomb  calorimeter 
(p.  174).  The  rise  in  temperature  of  a  known  weight  of  sur- 
rounding water  gives  the  number  of  calories.  The  coal  is  set  on 
fire  by  a  wire  heated  electrically.  Engineers  use  the  number  of 
British  Thermal  Units  (1  B.T.U.  =  heat  required  to  raise  1  Ib. 
of  water  1°  F.)  developed  by  1  pound  of  coal.  The  number  of 
B.T.U.  =  1.8  X  number  of  calories  per  gram  of  coal.  Bitumi- 
nous coals  give  much,  and  widely  varying  amounts  of  volatile 
matter,  while  anthracite  gives  very  little.  The  ash  is  the  mineral 


410 


COLLEGE    CHEMISTRY 


Water. 

Vola- 
tile 
matter. 

Fiied 
car- 
bon. 

Ask. 

Sul- 
phur. 

Cal.  per 
lg- 

Wood    

20  0 

49  0 

30  0 

1  0 

3  100 

Peat  

20  0 

51  6 

25  0 

3  2 

0  2 

4  270 

Bituminous    

1  3 

36  7 

53  5 

8  5 

1  7 

7,800 

Semi-bituminous  

4.0 

16.0 

68.5 

11 

0  5 

7,510 

Anthracite  

3  0 

5.6 

80.5 

10.9 

0.8 

8,000 

Charcoal 

3  2 

4  2 

90  7 

1  7 

7  580 

Coke 

2  5 

1  3 

86  3 

12  4 

1  3 

7  770 

Petroleum  .       .       .   . 

11  000 

matter  of  the  original  plants,  with  rock  material  in  many  speci- 
mens. For  coal  gas,  and  even  for  coke,  a  coal  high  in  volatile 
matter  is  chosen.  For  water  gas  (p.  386)  anthracite  or  coke  is 
employed. 

If  the  heat  of  combustion  of  a  coal  is  known,  the  amount  of 
steam  it  should  furnish  can  be  calculated.  It  takes  100  cal.  to 
raise  1  g.  of  water  from  0°  C.  to  100°  C.  and  540  cal.  more  to 
convert  it  into  steam.  If  the  quantity  of  steam  is  too  small,  the 
furnace,  draft,  or  firing  is  defective.  Too  much  draft,  for  ex- 
ample, merely  adds  additional,  useless  air  to  be  heated.  If  the 
flue  gas,  when  analyzed,  contains  only  3  per  cent  of  carbon  dioxide, 
instead  of  the  normal  12  per  cent,  then  for  every  ton  of  coal 


FIG.  110. 


burned,  52  tons  of  unnecessary  air  were  raised  to  the  temperature 
of  the  furnace.  Tests  of  this  kind  can  control  the  efficiency  of 
every  device  in  a  modern  factory,  and  they  ought  to  be  in  uni- 
versal use. 


Coal  Gas. —  The  gas  plant  (Fig.  110)  includes:    (1)  The  fire- 
brick retorts  in  which  the  coal  is  heated  to  1300°,  (2)  the  hydraulic 


THE    CAKBOHYD RATES  AND  RELATED   SUBSTANCES      411 

main,  a  wide  iron  pipe  above  them  in  which  the  tar  collects, 

(3)  the  condenser  and  wash  box  for  cooling  and  condensing  oils, 

(4)  the  scrubbers  where  the  ammonia  is  taken  out  by  water, 

(5)  the  purifier  where  hydrogen  sulphide  is  absorbed  by  hydrated 
ferric  oxide  and  (6)  the  holder  where  the  gas  collects. 

One  short  ton  (2000  Ibs.)  of  the  bituminous  coal  in  the  above 
table  gave:  Gas  10,500  cu.  ft.  with  13  candle  power,  coke  1325  Ibs., 
ammonia  5  Ibs.  (  =  20  Ibs.  (NH^SO*  worth  $60  per  ton),  and  tar 
12  gallons.  The  components  of  the  gas  were:  Illuminants  3.8, 
heating  gases  90.2,  impurities  6.0.  Calorific  power  of  gas  610 
B.T.U.  per  cu.  ft.;  sp.  gr.  (air  =  1)  0.43. 

The  tar  is  frequently  distilled  fractionally  and  yields:  Benzene 
CeH6,  from  which  aniline  and  many  dyes  and  drugs  are  prepared; 
naphthalene  Ci0H8,  sold  as  moth-balls,  and  the  starting  point  for 
synthetic  indigo;  anthracene  CuHio,  from  which  valuable  dyes 
such  as  alizarin  and  indanthrene  are  made;  phenol  or  carbolic 
acid  (p.  349),  and  other  useful  substances.  A  rougher  separation 
yields  tar  and  pitch,  for  road-making,  preserving  timber,  and 
waterproofing  roofs. 

Coke.  —  The  beehive  coke  oven  is  a  brick  structure  shaped  like 
a  beehive,  with  an  additional  opening  at  the  top.  The  coal  which 
fills  it  burns  with  a  limited  supply  of  air.  All  the  vapors  and  gas 
burn  at  the  upper  opening,  and  the  ammonia  and  tar  and  com- 
bustible gas  are  therefore  wasted  (cf.  p.  340). 

The  by-product  coke  oven  is  a  good  deal  like  a  gas  plant.  The 
chief  difference  is  that  the  heating  is  arranged  so  as  to  decompose 
as  much  of  the  volatile  matter  as  possible,  and  cause  it  to  leave  its 
carbon  in  the  retort.  The  gas  is  therefore  poor  in  illuminants,  but 
excellent  as  fuel.  The  ammonia  and  tar  are  diminished  in  amount, 
but  still  valuable  products.  The  yield  of  coke  is  about  73  per 
cent  of  the  original  coal,  against  66  per  cent  from  the  beehive 
oven. 

Burning  coke  gives  a  higher  temperature  than  does  coal,  be- 
cause no  heat  is  used  in  vaporizing  moisture  and  volatile  matter. 
For  the  same  reason,  it  burns  without  flame.  Because  of  these 
and  other  properties,  it  is  employed  in  immense  quantities  in  re- 
ducing ores  of  iron  in  the  blast  furnace,  as  well  as  for  many  other 
purposes. 


CHAPTER  XXXI 
ORGANIC  ACIDS  AND  SALTS.     ALCOHOLS,  ESTERS.     FOODS 

THUS  far,  one  acid,  acetic  acid,  and  two  alcohols,  methyl  and 
ethyl  alcohol,  have  been  mentioned.  But  there  are  whole  series 
of  organic  acids,  corresponding  to  the  series  of  hydrocarbons. 

Organic  Acids  and  Their  Salts.  —  The  general  formula  of  the 
saturated  series  of  monobasic  acids  is  H(CO2CnH2n+i).  Thus: 

Formic  acid  (n  =  0),  HCCCfeH).  Palmitic  acid  (n  =  15),  H(CO2Ci6H3i). 
Acetic  acid  (n  =  1),  HCCOzCHj).  Stearic  acid  (n  =  17),  H(CO2Ci7H36). 
Butyric  acid  (n  =  2),  H(COjC,H6). 

Formic  (Lat.,  an  ant)  acid  is  secreted  by  red  ants,  and  is  found 
in  stinging  nettles.  Formic  (b.-p.  100.1°),  acetic,  and  butyric 
acids  are  liquids.  Palmitic  and  stearic  acids  are  solids,  and  are 
mixed  with  paraffin  in  making  candles. 

Acids  containing  relatively  less  hydrogen  are  unsaturated.  Thus, 
oleic  acid  (n  =  17)  is  H(C02Ci7H33). 

The  acids  with  large  molecular  weight  are  insoluble  in  water. 
All  the  acids,  however,  react  with  sodium  hydroxide  solution, 
giving  soluble  salts.  Thus,  palmitic  acid  gives  sodium  palmitate : 

NaOH  +  H(CO2Ci5H3i)  <=±H2O  +  Na(C02Ci5H3i). 

Other  salts  are  sodium  formate  (p.  385)  Na(C02H),  sodium  acetate 
Na(C02CH3),  sodium  stearate  Na(C02Ci7H35),  sodium  oleate 
Na(C02Ci7H33).  Common  soap  is  a  mixture  of  the  last  two  salts 
with  sodium  palmitate. 

Later,  in  discussing  fats  and  soap,  it  will  be  convenient  to  abbre- 
viate the  formulae.  A  monobasic  acid  will  be  indicated  by  the 
formula  HCO2R  and  a  salt  by  NaC02R  or  Ca(CO2R)2,  where  R 
stands  for  a  hydrocarbon  radical  or  group  of  atoms,  such  as 
CnH2n+i.  In  organic  chemistry  a  radical  is  not  always  able  to 
form  an  ion.  Here  the  ion  is  C02R~. 

412 


ALCOHOLS,    ESTERS 


413 


There  are  also  dibasic  acids.  Oxalic  acid  (p.  385)  H2C204  is  the 
simplest  of  these.  It  may  be  made  by  the  oxidation  of  sugar  with 
nitric  acid.  The  white  crystals  used  in  the  laboratory  are  the 
hydrate  H2C204,2H20. 

Alcohols.  —  Methyl  alcohol  CH3OH  (p.  408)  and  ethyl  alco- 
hol C2H5OH  (p.  406)  are  the  first  two  members  of  the  series 
C«H2n+iOH.  There  are  also  many  alcohols  with  more  than  one 
OH  group  in  each  molecule.  Of  these,  the  one  we  shall  presently 
encounter  is  glycerine  C3H5(OH)3.  The  sugars  are  alcohols  with 
several  hydroxyl  radicals. 

Esters.  —  When  an  acid  and  an  alcohol  are  mixed,  an  ester 
and  water  are  formed.  The  action  is  slow  and,  being  reversible, 
is  always  incomplete.  But,  by  introduction  of  a  dehydrating 
agent  like  concentrated  sulphuric  acid,  the  water  is  removed  and 
the  change  brought  to  completion.  Thus,  ethyl  alcohol  and  acetic 
acid,  when  warmed  with  sulphuric  acid,  give  ethyl  acetate: 
C2H5OH  +  HC02CH3  <=±  C2H5C02CH3  +  H20. 

This  action  has  the  appearance  of  a  neutralization  (p.  256),  but  is 
different  in  several  ways.  Alcohol  is  not  a  base,  and  in  aqueous 
solution  it  does  not  conduct  electricity.  Then,  true  neutralization 
takes  place  instantly,  while  the  foregoing  action,  and  all  like  it, 
proceed  very  slowly.  Thus,  although  acetic  acid  is  a  true  acid, 
it  is  not  here  interacting  with  a  base.  f  ;3 

With  the  assistance  of  a  dehydrating  agent,  similar  actions  take 
place  between  any  alcohol  and  any  acid  (organic  or  inorganic). 
The  action  of  cellulose  or  of  glycerine  with  nitric  acid  (p.  3 
such  an  interaction.     Again,  for  example: 

C3H5(OH)3  +  3H(C02CH3)  ^  C3H5(C02CH3)3  +  3H20, 

glycerine  acetic  acid  glyceryl  acetate 

C3H5(OH)3  +  3H(C02C17H35)  ^  C3H5(C02C17H35)3  +  3H20. 
glycerine  stearic  acid  glyceryl  stearate 

The  glyceryl  radical  C3H5m  is  trivalent,  and  takes  the  place  of 
three  atoms  of  hydrogen.     The  products  are  named  as 
were  salts,  but  they  are  not  ionized  in  solution  and  do  n 
with  acids,  bases,  salts,  by  instantaneous  double  decomposition, 


414  COLLEGE    CHEMISTRY 

as  do  true  salts.  To  distinguish  them  from  salts,  they  are  called 
esters  —  R'(C02R). 

Fats  and  Animal  and  Vegetable  Oils.  —  The  fats  and  oils 
found  in  animal  tissue,  or  pressed  from  seeds  of  plants,  are  com- 
posed mainly  of  esters.  Beef  fat  is  a  mixture  of  about  three- 
fourths  glyceryl  palmitate  (palmitin)  C3H6(CO2Ci5H3i)3  and 
glyceryl  stearate  (stearin)  C3H5(CO2Ci7H35)3,  along  with  one- 
fourth  glyceryl  oleate  (olein)  C3H6(CO2Ci7H33)3.  Lard  (hog  fat) 
contains  a  much  larger  proportion  of  the  last  (60  per  cent)  and  is 
therefore  softer.  Butter  contains  the  same  esters,  along  with  some 
water  and  some  glyceryl  butyrate  (butyrin)  C3H6(CO2C3H7)3. 
Olive  oil  contains  much  olein  (75  per  cent).  Cottonseed  oil  is 
similar  in  composition,  and  is  used  as  a  substitute  for  olive  oil,  or 
for  butter  in  cooking. 

All  these  fats  and  oils  contain  a  certain  proportion  of  the  free 
organic  adds  (see  p.  420).  These  oils  must  not  be  confused  with 
mineral  oils,  which  are  mixtures  of  hydrocarbons. 

As  regards  physical  properties,  these  oils  are  all  insoluble  in  water, 
and  the  heavier  ones  also  in  cold  alcohol.  They  dissolve  readily, 
however,  in  ether,  benzene,  carbon  disulphide,  and  carbon  tetra- 
chloride.  Hence,  benzene  is  used  in  dry  cleaning  clothing  made 
of  silk  or  wool.  The  two  last  solvents  are  used  in  extracting  vege- 
table oils. 

Chemical  Properties  of  Fats  and  Oils.  —  All  fats  and  oils, 
when  boiled  with  water,  and  more  rapidly  when  heated  (200°) 
with  water  in  a  closed  vessel,  are  decomposed.  The  ester  is 
hydrolyzed,  and  the  actions  in  the  three  equations  last  given 
(p.  413)  are  reversed.  Thus,  with  stearin: 

C3H6(C02C17H35)3  +  3H20  ->  C3H5(OH)3  +  3HCO2C17H35, 
stearin  glycerine  stearic  acid 

and  when  the  mixture  is  cooled,  the  acid,  being  insoluble  in  water, 
forms  a  solid  cake  while  the  glycerine  is  in  solution  in  the  water. 
If  a  mixture  like  beef  fat  is  heated  with  water  in  this  way,  a  mix- 
ture of  palmitic,  stearic,  and  oleic  acids  is  obtained.  The  oleic 
acid  (liquid)  is  pressed  out,  and  the  residue  is  mixed  with  paraffin 
to  make  candles.  The  glycerine  is  separated  from  the  water  and 


ALCOHOLS,    ESTERS 


415 


used  in  making  nitroglycerine  (glyceryl  nitrate,  an  ester)  and  in 

medicine. 

When  the  fat  is  heated  with  aqueous  sodium  hydroxide,  the 
soluble  sodium  salts  of  the  acids  are  formed.  Since  these  salts  are 
known  as  soaps,  this  action  of  a  base  on  an  ester  is  called  saponi- 
fication: 

3NaOH-»C3H5(OH)3  +  3Na(CO2C17H35). 


stearin  glycerine  sodium  stearate 

(a  soap) 

When  common  salt  is  added  to  the  solution  ("salting  out"),  the 
sodium  salts  of  the  three  acids  (the  soap)  are  coagulated  and 
separated  as  a  floating  layer,  which  solidifies  when  cold. 
glycerine  is  contained  in  the  salt  solution. 

3  With  potassium  hydroxide,  the  potassium  salts  are  obtained, 
and  constitute  soft  soap. 

The  soaps  are  purified  by  redissolving  and  again  salting  out. 
Dyes  and  perfumes  are  often  added.  Floating  varieties  are  made 
by  beating  the  soap  before  it  solidifies,  and  so  introducing  bubbles 
of  air.  Fine  sand  or  pumice  is  added  to  make  scouring  soaps. 
Mixing  with  glycerine  or  sugar  gives  transparent  soap. 

Chemical  Properties  of  Soaps.  —  Since  the  soaps  are  soluble 
baits  of  sodium,  they  are  largely  ionized  in  solution  and  interact 
with  acids  by  double  decomposition: 

Na(C02C17H35)  +  HC1  ->  NaCl  +  H(CO2C17H35)  |  . 

The  acids,  being  insoluble,  are  precipitated.  They  also  enter  into 
double  decomposition  with  other  salts.  Thus,  hard  water,  con- 
taining compounds  of  calcium  and  magnesium  in  solution,  give 
precipitates  of  the  corresponding  salts.  For  example: 

2Na(C02C17H35)  +  CaSO4  -»  Na2SO4  +  Ca(C02C17H35)2l. 
Thus,  with  hard  water,  much  soap  is  wasted  in  precipitating  the 
''hardness." 

Colloidal  Suspension.  -To  explain  the  cleansing  power  of 
soap,  it  is  necessary  to  learn  more  about  colloids,  for  soap  11 
tion  is  essentially  colloidal. 


416 


COLLEGE    CHEMISTRY 


The  simplest  colloidal  suspensions  are  those  of  metals  like  gold 
and  platinum.  They  can  be  made  by  forming  an  electric  arc 
between  the  points  of  two  wires,  while  the  points  are  immersed  in 
water.  Liquids  of  various  colors,  depending  on  the  degree  of 
dispersion  (fineness  of  the  particles)  of  the  metal,  are  thus  formed 
Such  a  liquid  (1)  leaves  no  deposit  on  filter  paper,  (2)  shows  no 
elevation  in  the  boiling-point  of  the  solvent  and  (3)  no  depression 
in  the  freezing-point.  (4)  The  suspended  body  has  little  or  no 
tendency  to  diffuse  into  a  layer  of  the  pure  solvent.  In  conse- 
quence, if  the  colloidal  solution  is  placed  in  a  diffusion-shell,  which 
is  a  test-tube  shaped  tube  of  filter-paper  or  parchment,  immersed 
in  water,  none  of  the  colloid  escapes  through  the  pores  of  the  shell. 
Ordinary  solutes  escape  more  or  less  quickly,  according  to  their 
molecular  weight.  Hence,  a  diffusion  shell  can  be  used  to  separate 
a  mixture  of  colloidal  and  non-colloidal  material.  Thus,  salt,  if 
present  with  colloidal  starch,  or  sugar  if  present  with  colloidal 
gold,  can  be  removed  by  changing  the  water  round  the  shell  until 
no  more  salt  (or  sugar)  is  found  to  come  out.  This  process  is 
called  dialysis,  and  was  devised  by  Graham. 

(5)    The  most  striking  property  of  colloids  is   shown  by  the 
ultramicroscope.      In  a  perfectly  darkened  room,  a  converging 
beam    of    strong    light    is    sent    horizontally 
through  the  liquid   (Fig.   Ill)  and  the  place 
where  the  light  is  focussed  is   viewed  from 
above,  through  a  microscope.       Under  such 
circumstances,   a  true  solution  remains   per- 
fectly dark,  but  a  colloidal  suspension  shows 
^  minute  points  of  light,  first  studied  by  Tyn- 

in.  dall.  ^  Colloidal  gold,  solutions  of  soap,  starch, 

gelatine  and  dyes,  and  many  other  liquids 
exhibit  the  phenomenon.  The  points  of  light,  due  to  particles 
which,  although  minute,  contain  many  molecules,  show  also  a 
trembling  or  vibrating  movement,  first  noticed  by  a  botanist 
Brown  (1827)  and  called  the  Brownian  Movement.  The  motion 
is^due  to  collisions  of  the  moving  molecules  of  the  solvent 
with  the  suspended  particles  of  the  colloid  and,  when  the  sus- 
pension is  very  fine  (highly  "disperse"),  the  particles  shoot 
about  rapidly. 

Other  properties  of  colloidal  suspensions  are  discussed  below. 


ALCOHOLS,    ESTERS  417 

Theory    of    Colloidal    Suspension    and    Coagulation.— 

When  wires  from  a  battery  are  immersed  in  the  liquid,  the  par- 
ticles of  a  colloid  are  found  to  move  slowly  either  with  or  against 
the  positive  current.  The  phenomenon  is  called  electrophoresis. 
Apparently,  the  colloidal  particles  are  aggregates  of  molecules  of 
an  insoluble  substance,  collected  round  an  ion.  The  particles, 
although  relatively  large  in  proportion  to  the  charge,  move 
almost  as  rapidly  as  in  ionic  migration  (p.  231).  This  affords  an 
explanation  of  the  fact  that  the  particles  remain  suspended,  and 
do  not  settle.  They  are  individually  so  small  that  they  are  kept 
in  motion  by  collisions  with  the  molecules  of  the  solvent.  If  they 
could  unite  into  large  aggregates  —  like  the  particles  of  a  precipi- 
tate —  they  would  separate  like  any  ordinary,  insoluble  substance. 
But,  having  like  electrical  charges,  they  repel  one  another,  and  so 
remain  separate  and  in  suspension. 

Now  those  colloids  which  have  distinct  electrical  charges  can 
be  coagulated  or  flocculated,  and  so  precipitated  in  the  liquid,  by 
adding  a  solution  of  an  ionized  substance.  Thus,  colloidal  gold 
and  other  metals  are  negative,  and  an  equivalent  amount  of  a 
positive  ion,  usually  H+,  is  present  also.  When  a  salt  is  added, 
the  positive  ion  of  the  salt  attaches  itself  to  the  negative  colloidal 
metallic  particles,  neutral  bodies  result,  coagulation  can  now  occur, 
and  precipitation  follows.  Bivalent  ions  are  more  effective  than 
univalent  ones  (see  arsenic  trisulphide) .  Conversely,  a  positive 
colloid,  like  ferric  hydroxide,  is  coagulated  by  the  negative  ion  of 
the  salt,  and  more  easily  the  higher  the  valence  of  the  negative  ion. 
Furthermore,  one  colloid  will  coagulate  another  of  opposite  charge. 
Thus,  metaphosphoric  acid  is  a  negative  colloid  when  in  solution, 
while  ortho-  and  pyrophosphoric  acid  are  not  colloidal.  Albumin 
is  usually  a  positive  colloid.  Hence  (p.  372),  metaphosphoric  acid 
and  albumin  coagulate  and  precipitate  one  another,  while  the  other 
two  acids  have  no  action  on  albumin. 

Starch  and  gelatine  are  neutral  colloids,  and  are  not  easily 
coagulated. 

Soap  Solution  Colloidal,  Salting  Out.  —  Soap  solution,  un- 
der the  ultramicroscope,  is  seen  to  contain  suspended  particles. 
A  test  with  litmus  also  shows  that  the  soap  is  partly  hydrolyzed: 
Na(C02R)  +  H20<=±H(C02R)  +  NaOH. 


418  COLLEGE   CHEMISTRY 

Being  a  salt  of  a  little  ionized  acid,  the  negative  ion  of  the  soap 
tends  to  combine  with  the  H+  of  the  water:  H+  +  (C02R)~  <=± 
H(C02R),  leaving  the  ions  of  sodium  hydroxide.  Now  the  acid 
thus  set  free  combines  with  the  undissociated  molecules  of  the  salt 
to  form  an  acid  salt  NaH(CO2R)2-  This  acid  salt  is  insoluble,  but 
remains  in  colloidal  suspension  as  a  negative  colloid.  When  a 
strong  solution  of  common  salt  (or  even  excess  of  sodium  hydrox- 
ide) is  added,  the  positive  ion  Na+  is  adsorbed  by  the  negative 
colloid  (the  acid  salt)  and  the  latter  is  coagulated.  In  coming  out 
as  a  precipitate,  it  seems  to  adsorb  most  of  the  sodium  hydroxide, 
so  that  the  precipitate  has  the  composition  of  soap. 

The  Cleansing  Power  of  Soap.  Emulsions.  —  As  a  cleanser, 
soap  solution  —  or  suspension,  as  we  should  now  call  it  —  has  two 
properties.  It  removes  oil  and  grease  (insoluble  liquids)  by 
forming  an  emulsion  with  them.  It  also  removes  minute  solid 
particles  of  dirt,  by  taking  the  dirt  into  suspension  (next  section) . 

When  an  oil,  such  as  kerosene,  is  violently  shaken  with  water, 
both  liquids  are  broken  into  minute  droplets,  and  an  opaque  mix- 
ture results.  The  droplets  of  each  liquid,  however,  quickly  join 
together  and  soon  the  mass  clears  up  and  shows  the  two  liquids  in 
separate  layers.  If,  however,  a  colloidal  suspension  is  used, 
instead  of  pure  water,  the  droplets  unite  much  more  slowly,  if  at 
all,  and  a  more  or  less  permanent,  opaque,  rather  viscous  mass 
results.  Such  a  mixture  of  two  mutually  insoluble  liquids  is  called 
an  emulsion.  Thus,  a  few  drops  of  soap  solution  will  cause  the 
kerosene  and  water  to  remain  much  longer  in  the  condition  of  an 
emulsion.  Similarly,  vinegar  and  olive  oil,  when  vigorously 
beaten  (French  dressing)  separate  rather  quickly  into  two  layers. 
But  if  the  yolk  of  an  egg  (colloidal)  is  added  to  the  vinegar,  a  stiff, 
almost  solid  mass  can  be  made  (Mayonnaise  dressing)  which  will 
remain  permanently  emulsified.  In  removing  grease,  therefore, 
rubbing  with  soap  solution  turns  the  grease  into  suspended  drop- 
lets (emulsifies  it),  and  so  the  grease  can  be  washed  away. 

This  behavior  of  a  colloid  can  be  explained.  When  the  kerosene 
and  water  are  divided  into  droplets,  with  a  great  increase  in  the 
total  surface,  and  in  the  surface  energy  of  both,  the  surface  tension 
of  water,  which  is  great,  favors  the  reunion  of  the  drops,  with 
diminishing  surface,  and  dissipation  of  the  surface  energy.  Now, 


ALCOHOLS,   ESTERS  419 

while  ordinary,  dilute  solutions  have  a  surface  tension  close  to  that 
of  water,  colloidal  solutions  (such  as  0.5  per  cent  soap  solution) 
have  a  very  low  surface  tension.  Hence,  the  tendency  to  di- 
minish the  surface  of  droplets  of  soap  solution,  by  coalescence,  is 
slight  and  ineffective.  Furthermore,  as  predicted  by  Willard 
Gibbs  of  Yale  University,  and  proved  by  experiment,  a  colloid 
has  the  peculiarity  that  it  tends  to  reach  a  higher  concentration 
in  the  surface  layer  than  it  has  elsewhere  in  the  liquid.  When  the 
colloid  has  adjusted  itself  to  a  state  of  equilibrium,  in  this  regard, 
it  resists  a  decrease  in  the  surface  (which  would  increase  its  concen- 
tration beyond  the  equilibrium  value),  just  as  much  as  it  resists  an 
increase,  which  would  diminish  its  concentration.  The  emulsion 
of  a  colloidal  suspension  with  an  immiscible  liquid  is  thus  a  stable 
condition. 

Experiments  confirming  this  view  are  easily  made.  If  a  solu- 
tion of  a  dye  like  methyl  violet  (colloid)  be  shaken  violently,  and 
the  froth  (large  surface  in  proportion  to  quantity  of  liquid)  be 
separated,  it  is  found  that  the  liquid  produced  by  the  subsidence  of 
the  froth  (an  emulsion  with  air  is  not  permanent),  is  darker  in 
shade,  and  contains  more  dye,  than  an  equal  amount  of  the  original 
liquid.  Soap  solution,  after  being  shaken  likewise,  contains 
relatively  more  soap  in  the  froth  than  in  the  liquid. 

Adsorption  of  Colloidal  Matter.  —As  we  have  seen  (p.  409), 
when  liquids  containing  colloidal  substances,  such  as  dyes  and 
natural  organic  coloring  matters,  are  shaken  with  pulverized 
charcoal,  the  colloid  is  adsorbed  by  the  charcoal  —  that  is,  it  ad- 
heres to  the  surface  of  the  particles  of  the  charcoal.  This  principle 
is  used  in  decolorizing  sugar  (p.  405)  and  in  "  bleaching  "oils. 
Now,  soap  is  also  removed  by  shaking  with  charcoal  or  with  u 
fusorial  earth,  in  the  same  fashion. 

Pulverized  charcoal  is,  relatively,  a  coarse  powder, 
which  is  very  finely  divided  carbon,  be  freed  from  oil  or  grease  by 
washing  with  ether,  it  gives  a  loose,  non-caking  powder, 
powder  be  shaken  with  water,  it  settles.    If  it  be  shaken  wit! 
dilute  soap  solution,  it  remains  in  suspension,  and  the  liquid  re- 
sembles ink.     The  particles  are  so  fine  that,  instead  of  carrying 
down  the  colloidal  soap,  and  forming  a  precipitate,  as  char 
does,  they  attach  themselves  to  the  colloidal  soap,  and  re^^M 


420  COLLEGE    CHEMISTRY 

suspended.  This  is  therefore  adsorption,  with  the  difference  from 
the  ordinary  phenomenon,  that  the  colloid  carries  off  the  adsorb- 
ent, instead  of  the  adsorbent  carrying  down  the  colloid.  Now 
dirt  is  composed  largely  of  soot,  and  equally  fine  particles  of  other 
substances.  Hence,  the  soap  first  emulsifies  the  oil  on  the  hands 
(or  on  soiled  linen,  for  example)  and  then  adsorbs  the  particles  of 
dirt  which  are  thus  set  free. 

Formerly,  soap  solution  was  supposed  to  remove  grease  (and 
soot?)  because  of  its  slight  alkaline  reaction,  due  to  hydrolysis. 
This  explanation  must  be  given  up,  because:  (1)  an  alkali  so 
dilute  that  it  exists  in  equilibrium  with  the  free  fatty  acid,  can  not 
possibly  saponify  the  ester  contained  in  a  grease  spot.  (2)  Pure 
alkali  of  the  same  concentration  (or  stronger)  has  no  more  emulsi- 
fying power  than  has  water.  Such  an  alkaline  solution  will 
indeed  emulsify  an  animal  or  vegetable  oil  (cod-liver  oil,  cotton 
oil,  castor  oil),  but  it  does  so  by  interacting  with  the  free  fatty  acid 
always  present  in  such  oil  (p.  414)  and  forming  therefrom  a  soap. 
Such  an  alkaline  solution  does  not  emulsify  kerosene,  although 
soap  solution  does.  The  emulsifying  agency  in  each  case  is  a  soap. 
(3)  Very  dilute  alkali  has  no  more  effect  upon  soot  than  has  water, 
—  but  soap  solution  takes  clean  (greaseless)  soot  instantly  into 
permanent  suspension.  (4)  An  aqueous  solution  of  saponin 
C32H540i8,  obtainable  from  several  plants,  although  it  contains  no 
alkali,  lathers,  emulsifies,  and  adsorbs  dirt,  just  as  does  soap.  It 
is  a  colloid. 

CYANOGEN 

Cyanogen  C^N^.  —  This  compound  is  prepared  by  allowing  a 
solution  of  cupric  sulphate  to  trickle  into  a  warm  solution  of  potas- 
sium cyanide.  The  cupric  cyanide,  at  first  precipitated,  quickly 
decomposes,  giving  cuprous  cyanide  and  cyanogen: 

2KNC  +  CuS04  ->  Cu(NC)2i  +  K2S04, 
2Cu(NC)2  -»  2CuNC  +  C2N2| . 

Cyanogen  is  a  very  poisonous  gas  with  a  characteristic,  faint  odor. 

Hydrocyanic  Acid  HNC.  —  This  acid,  called  also  prussic  acid, 

is  most  easily  made  by  the  action  of  an  acid  upon  a  cyanide  (see 
Potassium  cyanide),  followed  by  distillation.  It  is  a  colorless  liquid 


FOODS  421 

boiling  at  26.5°.  It  has  an  odor  like  that  of  bitter  almonds,  and  is 
highly  poisonous.  In  aqueous  solution  it  is  an  extremely  feeble 
acid.  Hydrocyanic  acid  and  the  cyanides  are  unsaturated  com- 
pounds, a  fact  which  is  illustrated  in  the  two  following  paragraphs. 

Cyanates  and  Thiocyanates.  —  When  potassium  cyanide  is 
fused  and  stirred  with  an  easily  reducible  oxide,  like  lead  oxide 
(PbO),  the  metal  (for  example,  the  lead)  collects  at  the  bottom  of 
the  iron  crucible  in  molten  form,  and  potassium  cyanate  KNCO  is 
produced: 

KNC  +  PbO  ->  KNCO  +  Pb. 

Ammonium  cyanate  NH4NCO,  when  dissolved  in  warm  water, 
undergoes  a  profound  change.  It  turns  into  urea  (NH^CO  (car- 
bamide), which  has  the  same  composition.  The  former  is  a  salt, 
and  is  ionized,  and  enters  into  double  decompositions :  the  latter 
is  not  "ionized,  but  is  like  ammonia,  able  to  combine  with  acids. 
This  is  an  example  of  internal  rearrangement  (p.  148).  Two  sub- 
stances, of  the  same  composition  and  molecular  weight,  but 
different  chemical  behavior,  are  called  isomers.  Urea  is  one  of  the 
forms  in  which  nitrogen  is  eliminated  by  animals.  The  prepara- 
tion of  what  seemed  to  be  a  typical  product  of  life  from  a  substance 
(ammonium  cyanate)  easily  made,  if  necessary,  from  the  four 
elements  themselves,  was  the  first  case  of  the  artificial  production 
of  an  organic  compound,  and  created  a  great  sensation  when  it  was 
first  accomplished  by  Wohler  in  1828. 

When  potassium  cyanide  in  aqueous  solution  is  boiled  with 
sulphur  or  with  a  polysulphide  (p.  274),  it  is  converted  into  potas- 
sium thiocyanate  KCNS.  This  salt,  or  ammonium  thiocyanate 
NH4CNS,  is  used  in  testing  for  ferric  ions  on  account  of  the  deep- 
red  color  of  ferric  thiocyanate  (cf.  p.  182). 

FOODS 

Plants  and  animals  contain  substances  which  are  similar  in 
composition,  such  as  sucrose  and  lactose  (p.  404),  starch  and 
glycogen  (C6Hi0  05)2,  animal  fats  and  vegetable  oils  (both  esters). 
Albumins  and  other  proteins  are  found  in  both.  They  differ,  how- 
ever, markedly  in  the  sources  of  these  substances.  The  plant 
uses  simple  materials,  like  carbon  dioxide,  water,  and  potassium 


422 


COLLEGE    CHEMISTRY 


nitrate.    The  animal  can  make  no  use  of  these  substances  —  it 
must  be  fed  on  complex  compounds  of  animal  or  vegetable  origin. 

Foods.  —  Since  the  animal  is  continuously  eliminating  carbon 
dioxide,  moisture,  compounds  of  nitrogen,  salts,  and  other  sub- 
stances, and  is  also  giving  off  heat,  these  materials  must  be  re- 
placed, and  fuel  must  be  furnished.  Like  the  plant,  an  animal 
can  absorb  only  dissolved  material.  But  it  prepares  its  own 
solutions  in  a  remarkable  laboratory,  the  digestive  tract.  The 
production  of  suitable  soluble  substances  is  called  digestion. 

The  table  shows  the  chief  components  of  animal  food,  and  the 
proportion  in  which  each  is  present  in  the  chief  foods  used  by  man : 


Water. 

Protein. 

Fat. 

Carbo- 
hydrate. 

Ash. 

Beef  (lean)     •. 

73  8 

22   1 

2  9 

« 

1  2 

Cod 

82  6 

15  8 

0  4 

1  2 

Eggs 

73  7 

14  8 

10  5 

1  0 

Milk  *                 .... 

87  0 

3  3 

4  0 

5  0 

0  7 

Butter  

11  0 

1  0 

85  0 

3  0 

Cheese  (cheddar)      .... 
Oatmeal 

27.4 
7  3 

27.7 
16  1 

36.8 
7  2 

4.1 
67  5 

4.0 
1  9 

Wheat  flour 

11  9 

13  3 

1  5 

72  7 

0  6 

Beans  (dried)     

12  6 

22  5 

1  8 

59  6 

3  5 

Almonds  

4  8 

21  0 

54  9 

17  3 

2  0 

Maize  (green  corn)  .... 
Potatoes  
Lettuce    

75.4 
78.3 
94.7 

3.1 
2.2 
1.2 

1.1 
0.1 
0.3 

19.7 
18.4 
2.9 

0.7 
1.0 
0.9 

*  The  emulsified  fat  separates  slowly  as  the  cream;  the  protein  (casein,  colloidally  suspended 
in  the  skim  milk)  is  coagulated  by  rennet  and  constitutes  cheese;  the  carbohydrates  (lactose, 
a  sugar)  is  then  left  in  the  water,  along  with  inorganic  salts. 

We  observe  that  the  common  animal  foods,  except  milk,  containing 
lactose  (p.  404),  carry  no  carbohydrates  (the  ox  liver  contains 
about  2  per  cent  of  glycogen) ;  that  potatoes  and  corn,  when  dried, 
are  nearly  all  starch;  that  lean  meat,  dry,  is  all  protein;  that  some 
seeds  (wheat  and  beans)  contain  little  fat,  some  (oats)  much  more 
fat,  and  some  (almonds  and  nuts)  a  large  amount;  and  that 
lettuce  is  mainly  water,  with  useful  inorganic  salts  in  solution,  and 
cellulose. 

The  proteins,  several  of  which  have  been  mentioned  (pp.  312, 
340,  359)  are  white,  amorphous  substances  containing,  besides 


FOODS  423 

carbon,  hydrogen,  and  oxygen,  a  large  proportion  of  nitrogen 
(16  per  cent),  some  sulphur  (1  per  cent)  and  frequently  iron  and 
phosphorus  as  well. 

Digestion.  —  The  process  of  rendering  the  constituents  of  food 
soluble  is  like  fermentation  (p.  406)  —  it  is  performed  mainly  by 
enzymes.  Each  class  of  components  is  handled  by  one  or  more 
enzymes.  Thus,  starch  (in  bread  and  potatoes)  is  partly  digested 
during  mastication  by  ptyalin  (an  amylase,  p.  406)  in  the  saliva, 
and  partly  by  amylopsin  in  the  small  intestine.  The  resulting 
maltose  is  decomposed  into  glucose  by  another  enzyme,  and  passes 
into  the  blood  where  it  is  oxidized,  furnishing  heat.  In  diabetes, 
much  of  the  glucose  escapes  oxidation,  and  is  eliminated.  Again, 
the  fats  are  hydrolyzed  into  the  acids  and  glycerine  by  lipases 
(fat-splitting  enzymes)  in  the  bile,  and  the  acids  go  into  solution 
(probably  colloidal).  The  acids  "and  glycerine  recombine  to  form 
fats  in  the  blood,  and  are  either  deposited  in  the  tissues  or  oxidized. 
Finally,  the  proteins  are  changed  in  a  similar  way  into  peptones 
which  are  soluble  in  water,  and  in  this  form  are  able  to  pass  through 
the  wall  of  the  intestine. 

Fuel  Value  of  Food.  —  Although  food  is  required  to  replace 
waste,  much  of  it  is  needed  to  furnish  energy,  by  its  oxidation,  so 
that  muscular  movements  may  be  maintained,  and  the  tempera- 
ture of  the  body  kept  up  to  its  normal  value  (37°  C.).  Thus,  the 
fuel  values  of  foods  are  important.  The  average  fuel  values,  ex- 
pressed in  large  calories  (I  Cal.  =  1000  cal.  as  previously  denned, 
p.  174),  per  gram,  are: 

Carbohydrates,  4  Cal.     Fats,  9  Cal.     Proteins,  4  Cal. 

The  fuel  values  per  pound  (453.6  g.)  are  453.6  times  greater. 

Healthy  life  cannot  be  maintained  on  one  kind  of  food  —  a 
mixed  diet  is  necessary.  In  general,  it  is  held  that  100  g.  of 
proteins  (giving  400  CaL)  per  day,  and  a  sufficient  amount  of  other 
foods  to  give  a  total  fuel  value  of  2200  cal.  is  enough  for  a  person 
doing  no  physical  labor.  When  physical  labor  is  involved,  larger 
values,  up  to  3800  cal.  per  day,  are  necessary.  The  data  in  the 
table  (p.  422)  will  enable  one  to  calculate  the  fuel  value  of  100  g. 
(or  of  1  Ib.)  of  each  kind  of  food. 


424  COLLEGE    CHEMISTRY 

Exercises. —  1.   Make  the  graphic  formulae  of  ethyl  formate, 
ethylene  bromide  (p.  393),  ethyl  alcohol. 

2.  Make  equations  for  the  formation  of  palmitin  (p.  413),  the 
saponification  of  olein  (p.  415). 

3.  Prepare  a  summary  of  the  various  statements  that  have  been 
made  in  the  text  about  catalysis  (e.g.,  pp.  29,  59,  156,  160,  279, 
288,  341,  361,  406),  and  illustrate  fully. 

4.  Calculate  the  fuel  value  of  1  Ib.  each  of  (a)  oatmeal,  (6) 
potatoes,  (c)  lettuce. 

5.  Calculate  the  weights,  both  in  pounds  and  in  grams,  of  100 
Cal.  portions  of  (a)  eggs,  (6)  wheat  flour,  (c)  almonds,  (d)  lettuce. 

6.  At  current  market  prices,  what  would  be  the  cost  per  100 
Cal.  portion  of  beef,  cod,  butter,  and  wheat  flour,  respectively. 


CHAPTER  XXXII 
SILICON  AND   BORON 

IN  respect  to  chemical  relations  there  is  a  close  resemblance  be- 
tween silicon  and  carbon.  Silicon  gives  a  monoxide,  but  is  quad- 
rivalent in  all  its  other  compounds.  It  is  a  non-metallic  element. 

Occurrence.  —  Silicon,  unlike  carbon,  is  not  found  in  the  free 
condition.  In  combination  it  is  the  most  plentiful  element  after 
oxygen,  and  constitutes  more  than  one-quarter  of  the  crust  of  the 
earth.  The  oxide  is  silica  or  sand  (Si02),  and  this  oxide  and  its 
compounds  are  components  of  many  rocks.  In  the  inorganic 
world  silicon  is  the  characteristic  element  to  almost  as  great  an 
extent  as  is  carbon  in  the  organic  realm. 

Preparation  of  Silicon.  —  When  finely  powdered  magnesium 
and  sand  are  mixed,  and  one  part  of  the  mass  is  heated,  a  violent 
action  spreads  rapidly  through  the  whole: 

2Mg  +  Si02  ->  Si  +  2MgO. 

At  the  same  time,  and  especially  if  excess  of  the  metal  is  used, 
some  magnesium  silicide  Mg2Si  is  formed  also.  The  mixture  is 
treated  with  a  dilute  acid  which  decomposes  the  magnesium  oxide 
and  the  silicide  (see  below),  and  leaves  the  silicon  (amorphous) 
undissolved.  When  amorphous  silicon  is  dissolved  in  molten  zinc, 
the  mass,  when  solid,  contains  crystalline  silicon.  The  zinc  is 
removed  by  the  action  of  a  dilute  acid,  the  silicon  remaining  un- 
affected. 

Silicon  and  ferrosilicon  (an  alloy  of  iron  and  silicon)  are  now 
made  on  a  large  scale,  the  former  by  heating  sand  and  carbon,  the 
latter  by  heating  a  mixture  of  ferric  oxide  and  sand  with  carbon 
in  the  electric  furnace  (p.  377). 

Properties.  —  Amorphous  silicon  is  a  brown  powder.  It 
unites  with  fluorine  at  the  ordinary  temperature,  with  chlorine  at 

425 


426  COLLEGE    CHEMISTRY 

430°,  with  bromine  at  500°,  with  oxygen  at  400°,  with  sulphur  at 
600°,  with  nitrogen  at  about  1000°,  and  with  carbon  and  boron  at 
temperatures  attainable  only  in  the  electric  furnace.  It  is  dis- 
solved by  a  mixture  of  hydrofluoric  acid  and  nitric  acid,  giving 
silicon  tetrafluoride.  Crystallized  silicon  forms  black  needles  be- 
longing to  the  hexagonal  system.  Silicon  and  ferrosilicon  act 
readily  upon  a  cold  solution  of  sodium  hydroxide  (cf.  p.  56),  the 
ortho-  or  metasilicate  of  sodium  being  formed: 

Si  +  2NaOH  +  H20  ->  Na^SiOa  +  2H2  1  . 

This  is  one  of  the  sources  of  hydrogen  for  filling  balloons  and  air- 
ships. 

Silicon  Hydride  SiH±.  —  Silicon  differs  from  carbon  in  giving 
only  two  well-defined  compounds  with  hydrogen.  The  chief  one 
may  be  liberated  as  a  gas  by  the  action  of  hydrochloric  acid  upon 
magnesium  silicide: 

MgsSi  +  4HC1  ->  2MgCl2  +  SiH4. 

The  action  is  like  that  by  which  hydrogen  sulphide  is  made.  The 
gas  is  easily  inflammable,  and  burns  to  form  water  and  silicon 
dioxide.  When  heated  alone,  it  decomposes  into  its  constituents. 

Silicon  Tetrachloride  and  Tetrafluoride.  —  The  tetrachlo- 
ride  SiCU  is  formed  by  direct  union  of  the  free  elements.  It 
may  also  be  prepared  by  passing  chlorine  over  a  strongly  heated 
mixture  of  silicon  dioxide  and  carbon.  The  products  enter  a 
condenser  in  which  the  tetrachloride  assumes  the  liquid  form: 

2C12  +  SiO2  +  2C  ->  SiCU  +  2CO. 


Chlorine  is  unable  to  displace  oxygen  from  combination  with 
silicon,  and  has,  therefore,  when  alone,  no  effect  upon  sand.  In 
the  above  action,  therefore,  the  carbon  is  used  to  secure  the  oxygen 
while  the  chlorine  combines  with  the  silicon.  This  kind  of  inter- 
action was  formerly  used  for  making  many  chlorides  (e.g.,  BC13, 
A1C13,  CrCl3)  from  oxides,  before  simple  ways  of  obtaining  the 
elements  in  the  free  condition  were  known. 

Silicon  tetrachloride  is  a  colorless  liquid  (b.-p.  59°),  which  fumes 
strongly  in  moist  air,  and  with  water  precipitates  silicic  acid: 
SiCl4  +  4H20  ->  4HC1  +  Si(OH)4|. 


SILICON  427 

When  strong  hydrofluoric  acid  acts  upon  sand,  silicon  tetra- 
fluoride  SiF4  is  liberated: 

Si02  +  4HF  -» 2H20  +  SiF4. 

Since  the  water  interacts  with  the  tetrafluoride  (see  below),  the 
latter  is  usually  made  by  heating  sand  with  powdered  calcium 
fluoride  and  excess  of  sulphuric  acid.  In  this  way  the  hydrogen 
fluoride  is  generated  in  contact  with  the  sand,  and  at  the  same  time 
the  sulphuric  acid  renders  the  water  inactive.  Hydrofluoric  acid 
acts  in  a  corresponding  way  upon  all  silicates  (q.v.),  whether  these 
are  minerals  or  are  artificial  silicates  like  glass  (cf.  p.  206). 

Silicon  tetrafluoride  is  a  colorless  gas.  It  fumes  strongly  in 
moist  air,  and  acts  vigorously  upon  water.  This  interaction  is 
different  from  that  of  the  tetrachloride,  because  the  excess  of  the 
tetrafluoride  forms  a  complex  compound  with  the  hydrofluoric 
acid: 

SiF4  +  4H20  ->  Si(OH)4  (+  4HF).  (1) 

(4HF)  +  2SiF4  -»  2H2SiF6. (2) 

3SiF4  +  4H2O  -»  Si(OH)4  +  2H2SiF6. 

The  silicic  acid  is  precipitated  in  the  water,  and  may  be  separated 
by  filtration,  leaving  a  solution  of  hydrofluosilicic  acid. 

Hydrofluosilicic  Acid  H2SiFQ.  —  This  acid  is  stable  only  in 
solution.  When  the  water  is  removed  by  evaporation,  silicon  tet- 
rafluoride is  given  off,  while  most  of  the  hydrogen  fluoride  remains 
to  the  last.  Its  salts  are  decomposed  in  a  corresponding  way  when 
they  are  heated.  This  acid  is  used  in  analysis  chiefly  because  its 
potassium  salt  K2SiF6  is  one  of  the  few  salts  of  this  metal  which 
are  relatively  insoluble  in  water.  The  barium  salt  is  also  insoluble, 
but  most  of  the  salts  of  the  heavy  metals  are  soluble. 

Silicon  Dioxide  SiO2.  —  This  substance  is  found  in  many 
different  forms  in  nature.  In  large,  transparent,  six-sided  prisms 
with  pyramidal  ends  it  is  known  as  quartz  or  rock  crystal.  When 
colored  by  manganese  and  iron  it  is  called  amethyst,  when  by 
organic  matter,  smoky  quartz.  A  special  arrangement  of  the 
structure  gives  cat's  eye.  Amorphous  forms  of  the  same  material, 
often  colored  brown  or  red  with  ferric  oxide,  are  agate,  jasper,  and 
onyx,  the  last  much  used  in  making  cameos.  Slightly  hydrated 


428  COLLEGE    CHEMISTRY 

varieties  of  silica  are  the  opal  and  flint.  Forms  produced  by  or- 
ganisms are  sponges  and  infusorial  earth  (Tripoli).  The  latter 
is  used  in  scouring  materials  and  for  decolorizing  oils  (p.  419). 

Silica  is  found  in  the  hard  parts  of  straw,  of  some  species  of 
horsetail  (equisetum) ,  and  of  bamboo.  In  the  form  of  whetstones 
it  is  used  for  grinding.  The  clear  crystals  are  employed  in  making 
spectacles  and  optical  instruments  and  are  more  transparent  to 
ultra-violet  light  than  is  glass.  Pure  sand  is  used  in  glass  manu- 
facture (q.v.).  Recently,  small  pieces  of  chemical  apparatus  have 
been  manufactured  by  fusing  quartz  (m.-p.  1600°)  in  the  oxy- 
hydrogen  flame  or  the  electric  furnace.  The  material  does  not 
crystallize  on  cooling,  and  is  amorphous,  like  glass.  Owing  to  the 
low  coefficient  of  expansion  of  silica,  the  vessels  can  be  heated  red 
hot  and  chilled  in  cold  water  without  risk  of  fracture.  *  : 

Silicates.  Water  Glass.  —  Silicon  dioxide,  although  differing 
profoundly  from  carbon  dioxide  in  its  physical  nature,  nevertheless 
behaves  like  the  latter  chemically.  Thus,  when  boiled  with 
sodium  hydroxide  solution  it  forms  sodium  metasilicate  Na2SiOs 
or  orthosilicate  Na4Si04. 

Si02  +  2NaOH  -» Na^SiOs  +  H2O. 

The  salt  is  left  as  a  gelatinous  solid  (" soluble  glass")  when  the 
water  is  evaporated.  The  silicates  of  potassium  and  sodium  may 
also  be  obtained  by  boiling  sand  with  the  carbonates  of  these 
metals,  or,  more  rapidly,  usually  as  metasilicates  (see  below),  by 
fusing  the  mixture: 

Si02  +  K2C03  -»  K2Si03  +  C02T  . 

Water  glass  or  soluble  glass,  being  a  salt  of  a  feeble  acid  with  an 
active  base,  gives  a  solution  with  an  alkaline  reaction  (p.  271,  383). 
When  manufactured  for  commercial  use,  it  has  the  composition 
Na2Si205(Na2Si03,SiO2),  and  gives  a  less  alkaline  solution.  It  is 
used  as  a  filler  in  cheap  soaps,  for  fireproofing  and  waterproofing 
timber  and  textiles,  and  for  preserving  eggs. 

Silicic  Acid  H^SiO^.  —  When  acids  are  added  to  a  solution  of 
sodium  silicate,  silicic  acid  is  set  free.  After  a  little  delay  it  usually 
appears  as  a  gelatinous  precipitate.  When,  however,  the  silicate 


SILICON  429 

is  poured  into  excess  of  hydrochloric  acid,  no  precipitation  occurs. 
The  silicic  acid  remains  in  colloidal  suspension.  The  acid  is  ortho- 
silicic  acid: 

Na4Si04  +  4HC1  -»  4NaCl 
Na2Si03  +  2HC1  +  H20  -*  2NaCl 


but  the  gelatinous  precipitate,  when  dried,  loses  the  elements  of 
water.  There  seem  to  be  no  definite  stages,  indicating  the  exist- 
ence of  various  acids,  such  as  we  observe  with  phosphoric  acid. 
The  final  product  of  drying  is  the  dioxide.  Silicic  acid  is  a  very 
feeble  acid  and,  therefore,  gives  no  salt  with  ammonium  hydroxide 
(feeble  base). 

The  suspension  of  colloidal  silicic  acid  can  be  freed  from  the  acid 
and  sodium  chloride  (see  equation,  above)  by  dialysis  (p.  416). 
It  is  a  positive  or  a  negative  colloid,  according  to  the  mode  of 
preparation,  and  the  two  kinds  are  coagulated  by  addition  of  salts 
having  bivalent  negative  and  positive  ions,  respectively. 

Mineral  Silicates.  —  While  silicic  acid  is  the  ortho-acid 
ELiSiO4,  and  no  other  silicic  acids  have  been  made,  the  salts  are 
most  easily  classified  by  imagining  them  to  be  derived  from  various 
acids  representing  different  degrees  of  hydration  of  the  dioxide 
(cf.  p.  369),  or,  to  put  it  the  other  way,  different  degrees  of  dehy- 
dration of  the  ortho-acid.  The  following  equations  show  the 
relation  of  the  ortho-acid  to  some  of  the  silicic  acids  whose  salts 
are  most  commonly  found  amongst  minerals: 

H4SiO4-    H20-»H2SiO3  (=    H2O,SiO2)      Metasilicic  acid. 
2H4SiO4  -    H2O  -i  H6S207  (=  3H2O,2SiO2)  )  ,..  ...  .       .  , 

2H4Si04  -  3H20  ->  H2Si205  (=    H2O,2Si02)  I  DlSllwnc  aclds" 
3H4Si04  -  4H2O  ->  H4Si308  (=  2H20,3Si02)    Trisilicic  acid. 

Di-  and  trisilicates  are  those  derived  from  acids  containing  two  and 
three  units  of  silicic  anhydride,  respectively,  in  the  formula.  The 
valences  of  the  negative  radicals  of  the  acids  are  shown  by  the 
number  of  hydrogen  units  in  the  formulae. 

The  composition  of  minerals  is  often  exceedingly  complex.  This 
is  due  to  the  fact  that  amongst  them  mixed  salts  (p.  245)  are  very 
common,  in  which  the  hydrogen  of  the  imaginary  acid  is  displaced 
by  two  or  more  metals  in  such  a  way  that  the  total  quantity  of  the 


430  COLLEGE   CHEMISTRY 

metals  is  equivalent  to  the  hydrogen.  The  following  list  presents 
in  tabular  form  some  typical  or  common  minerals  arranged  accord- 
ing to  the  foregoing  classification  : 


ORTHOSILICATES    (I^SiC^)  METASILICATES    (H2Si03) 

Zircon,  ZrSi04  Wollastonite,  CaSi03 

Garnet,  Ca3Al2(Si04)3  Beryl,  Gl3Al2(Si03)6 

Mica,  KH2Al3(Si04)3  Talc  (soapstone),  H2Mg3(Si03)4 

Kaolin,  H2Al2(Si04)2,H2O  Asbestos,  Mg3Ca(SiO3)4 

DISILICATE    (H6Si207)  TRISILICATE    (H4Si3O8) 

Serpentine,  Mg3Si2O7,2H20     Orthoclase  (felspar),  KAlSi308 

It  will  be  seen  that  the  total  valence  of  the  metal  units  is  equal  to 
that  of  the  acid  radicals.  Thus,  in  beryl  there  are  six  equivalents 
of  glucinum  (beryllium)  and  six  of  aluminium,  taking  the  place  of 
twelve  units  of  hydrogen  in  (H2SiO3)6. 

Mica,  which  is  obtained  in  large  sheets  from  Farther  India,  is 
used  in  making  lamp  chimneys  and  as  an  insulator  in  electrical 
apparatus.  Kaolin,  or  clay,  like  mica,  is  an  acid  orthosilicate. 
Garnets  are  pulverized  in  manganese  steel  crushers  and  used  in 
making  sandpaper. 

Some  of  these  minerals  frequently  occur  mixed  together  as  regu- 
lar components  of  certain  igneous  rocks.  Thus,  granite  (p.  2) 
is  a  more  or  less  coarse  mixture  of  quartz,  mica,  and  felspar. 
Frequently  the  oblong,  flesh-colored  or  white  crystals  of  the  last 
are  very  conspicuous.  Sandstone  is  composed  of  sand  cemented 
together  by  clay  or  by  lime,  and  colored  brown  or  yellow  by  ferric 
oxide. 

The  high  melting-point  of  silica,  compared  with  carbon  dioxide, 
and  the  formation  of  these  complex  silicates,  indicate  that  the 
oxide  is  highly  associated  (Si02)«. 

BORON  B 

As  regards  chemical  relations,  boron,  being  a  uniformly  trivalent 
element,  is  a  member  of  the  aluminium  family  (see  Table  of  periodic 
system).  Yet  it  is  a  pronounced  non-metal,  and  its  oxide  and 
hydroxide  are  acidic:  aluminium  is  a  metal,  and  with  its  oxide 
and  hydroxide  basic  properties  predominate.  Boron  and  its  com- 
pounds really  resemble  carbon  and  silicon  and  their  compounds  in 
all  chemical  properties,  excepting  that  of  valence. 


BORON 


431 


Occurrence.  —  Like  silicon,  boron  is  found  in  oxygen  com- 
pounds, namely,  in  boric  acid  and  its  salts.  Of  the  latter,  sodium 
tetraborate  Na2B407,  or  borax,  came  first  from  India  under  the 
name  of  tincal.  It  constitutes  a  large  deposit  in  Borax  Lake  in 
California.  Colemanite,  Ca2B6Oii,5H20,  from  California,  fur- 
nishes a  large  part  of  the  commercial  supply  of  compounds  of 
boron. 

Preparation.  —  When  boric  oxide  is  heated  with  powdered 
magnesium  (B203  +  3Mg  ->  3MgO  +  2B),  black,  amorphous  boron 
can  be  separated  with  some  difficulty  from  the  borides  of  magnesium 
in  the  resulting  mixture.  When  excess  of  powdered  aluminium  is 
used,  hard  crystals  of  boron  are  formed. 

Properties.  —  Boron  unites  with  the  same  elements  as  does 
silicon  (p.  425),  but  with  somewhat  greater  activity.  Like  carbon 
(pp.  276,  355),  it  is  also  oxidized  by  hot,  concentrated  sulphuric  or 
nitric  acid,  the  product  being  boric  acid.  It  interacts  with  fused 
potassium  hydroxide,  giving  a  borate: 

2B  +  6KOH  -> 2K3B03  +  3H2. 

Boron,  when  heated  with  nitrogen,  unites  to  form  the  nitride  BN, 
a  white  solid.  When  heated  in  the  electric  furnace  with  carbon, 
it  forms  a  carbide  B6C.  This  substance  is  harder  than  carborun- 
dum, and  stands  next  to  the  diamond  in  this  respect  (Appendix  II). 

The  Halides  of  Boron.  —  By  combined  action  of  carbon  and 
chlorine  on  boric  oxide  (p.  426),  the  trichloride  of  boron  BC13  may 
be  made.  It  is  a  liquid  (b.-p.  18°)  which  fumes  strongly  in  moist 
air,  and  is  completely  hydrolyzed  by  water. 

Boron  trifluoride  BF3  is  made  by  the  interaction  of  calcium 
fluoride  and  sulphuric  acid  with  boron  trioxide.  The  mode  of 
preparation  and  the  properties  of  the  substance  recall  silicon 
tetrafluoride  (p.  427).  It  interacts  with  water,  like  the  latter, 
giving  boric  acid  and  hydrofluoboric  acid  HBF4: 

4BF3  +  3H20  ->  B(OH)3  +  3HBF4. 

Boric  Acid  and  Boron  Trioxide. —  "Boric  acid  (boracic  acid), 
H3B03  is  somewhat  volatile  with  steam,  and  is  found  in  Tuscany 


432  COLLEGE    CHEMISTRY 

in  jets  of  water  vapor  (soffioni)  which  issue  from  the  ground. 
Water,  retained  in  basins  of  brickwork,  is  placed  over  the  open- 
ings, and  from  this  water,  after  evaporation,  boric  acid  is  obtained 
in  crystalline  form.  As  boric  acid  is  a  very  feeble  acid,  and  withal 
little  soluble,  it  may  also  be  made  from  sulphuric  acid  and  con- 
centrated borax  solution.  It  crystallizes  on  cooling  the  mixture: 

NaAOy  +  H2S04  +  5H20  t?  Na2SO4  +  4H3BO3|. 

Boric  acid  crystallizes  from  water  in  thin  white  plates,  which  are 
unctuous  (like  graphite  and  talc)  to  the  touch.  Its  solubility  in 
water  is  4  parts  in  100  at  19°,  and  34  in  100  at  100°.  The  solution 
scarcely  affects  litmus.  The  green  tint  it  confers  on  the  Bunsen 
flame  is  used  as  a  test  for  the  acid.  At  100°  the  acid  slowly  loses 
water,  leaving  metaboric  acid  HBO2,  and  at  140°  tetraboric  acid 
is  formed :  4HB02  —  H2O  — »  H2B4O7.  Strong  heating  gives  the 
trioxide  B2O3,  a  glassy,  white  solid.  When  dissolved  in  water,  all 
these  dehydrated  compounds  revert  to  orthoboric  acid  HaBOs. 
The  solution  of  boric  acid  in  water  is  used  as  an  antiseptic  in 
medicine  (half-saturated,  2  per  cent  solution),  and  sometimes  as 
a  preservative  for  milk  and  other  foods. 

Borates.  —  Borates  derived  from  orthoboric  acid  are  practically 
unknown.  The  most  familiar  salt  is  borax  or  sodium  tetraborate. 
The  decahydrate  Na2B407,10H2O,  which  crystallizes  from  water  at 
27°  in  large,  transparent  prisms,  and  the  pentahydrate  which 
crystallizes  at  56°,  are  both  marketed.  They  are  made  by  crystal- 
lization of  native  borax.  In  Germany,  borax  is  prepared  from 
boracite,  found  at  Stassfurt,  by  decomposing  a  solution  of  the 
mineral  with  hydrochloric  acid: 

MgCl2,2Mg3B8015  +  12HC1  +  18H20  -»  7MgCl2  +  16H3B03. 

The  boric  acid  is  dissolved  in  boiling  water,  and  sodium  carbon- 
ate is  added:  4H3BO3  +  Na2C03  -»  Na2B407  +  6H2O  -f  CO2.  In 
California  it  is  made  from  colemanite  by  interaction  with  sodium 
carbonate. 

Since  boric  acid  is  a  feeble  acid,  borax  is  hydrolyzed  by  water, 
and  the  solution  has  a  marked  alkaline  reaction  (cf.  p.  271).  In 
a  0.1N  solution  (25°),  0.5  per  cent  is  hydrolyzed. 


BORON  433 

When  heated  with  oxides  of  metals,  sodium  tetraborate  behaves 
like  sodium  metaphosphate  (c/.  p.  371),  and  is  used  in  the  form  of 
beads  in  analysis.  If  its  formula  be  written  2NaBO2,B2Os  (c/.  p. 
369)  it  will  be  seen  that  a  considerable  excess  of  the  acid  anhydride 
is  contained  in  it,  and  that,  therefore,  a  mixed  metaborate  may  be 
formed  by  union  with  some  basic  oxide.  Thus,  with  a  trace  of 
cupric  oxide,  the  bead  is  tinged  with  blue,  from  the  presence  of 
a  compound  like  2NaBO2,Cu(B02)2.  In  welding  iron,  borax  is 
scattered  on  the  parts,  and  combines  with  the  oxide  to  form  a 
fusible  mixed  borate,  which  is  forced  out  by  the  pressure.  Borax 
is  also  mixed  with  glass  in  making  enamels  for  cooking  utensils. 

Exercises. —  1.  Compare  and  contrast  the  elements  carbon 
and  silicon,  and  their  corresponding  compounds. 

2.  What  would  be  the  interaction  between  aqueous  solutions  of 
an  ammonium  salt  and  of  sodium  orthosilicate  (c/.  p.  429)?  Why 
is  ammonium  silicate  completely  hydrolyzed  by  water? 


CHAPTER  XXXIII 
THE   BASE-FORMING  ELEMENTS 

IN  the  present  chapter  a  preliminary  view  of  the  chemistry  of 
the  metallic  elements  is  given. 

Physical  Properties  of  the  Metals.  —  Metals  show  what  is 
commonly  called  a  metallic  luster,  but,  as  a  rule,  they  do  so  only 
when  in  compact  form.  Magnesium  and  aluminium  exhibit  it 
when  powdered,  but  in  this  condition  most  metals  are  black. 

The  metals  can  all  be  obtained  in  crystallized  form,  when  a 
fused  mass  is  allowed  to  cool  slowly  and  the  unsolidified  portion 
is  poured  off.  In  almost  all  cases  the  crystals  belong  to  the  regu- 
lar system. 

The  metals  vary  in  specific  gravity  from  lithium,  which  is  little 
more  than  half  as  heavy  as  water  (sp.  gr.  0.59),  to  osmium,  whose 
specific  gravity  is  22.5.  Those  which  have  a  specific  gravity  less 
than  5,  namely,  potassium,  sodium,  calcium,  magnesium,  alumin- 
ium, and  barium,  are  called  the  light  metals,  and  the  others  the 
heavy  metals. 

Most  metals  are  malleable,  and  can  be  beaten  into  thin  sjieets 
without  loss  of  continuity.  Those  which  are  allied  to  the  non- 
metals,  however,  such  as  arsenic,  antimony,  and  bismuth,  are 
brittle.  The  order  of  the  elements  in  respect  to  this  property, 
beginning  with  the  most  malleable,  is :  Au,  Ag,  Cu,  Sn,  Pt,  Pb,  Zn, 
Fe,  Ni. 

The  tenacity  of  the  metals  places  them  in  an  order  different  from 
the  above.  It  is  measured  by  the  number  of  kilograms  which  a 
piece  of  the  metal  1  sq.  mm.  in  section  can  sustain  without  break- 
ing. The  values  are  as  follows:  Fe  62,  Cu  42,  Pt  34,  Ag  29,  Au  27, 
Al  20,  Zn  5,  Pb  2. 

The  hardness  is  measured  by  the  ease  with  which  the  material 
may  be  scratched  by  a  sharp,  hard  instrument.  Potassium  is  as 
soft  as  wax,  while  chromium  is  hard  enough  to  cut  glass  (Ap- 
pendix II). 

434 


THE   BASE-FORMING   ELEMENTS 


435 


The  temperature  at  which  the  metal  fuses  has  an  important 
bearing  on  its  manufacture.  Most  of  the  following  melting-points 
are  only  approximate: 


Mercury  .  . 
Potassium  .  . 
Sodium  .  .  . 
Tin  .  .  . 

-39° 
62° 
96° 
232° 

Zinc    .... 
Antimony 
Magnesium  . 
Aluminium   . 

419° 
630° 
651° 
659° 

Cast  iron  .  . 
Manganese  . 
Nickel  .  .  . 
Chromium  . 

1150° 
1260° 
1452° 
1520° 

Bismuth 

271° 

Silver      .    .    . 

960° 

Iron  (pure)  . 

1530° 

Cadmium  .  . 
Lead  .... 

321° 
327° 

Gold    .... 
Copper   .    .    . 

1063° 
1083° 

Platinum  .  . 
Tungsten  .  . 

1755° 
3267° 

It  will  be  seen  that  mercury  is  a  liquid,  that  potassium  and 
sodium  melt  below  the  boiling-point  of  water,  and  that  the  metals 
down  to  the  foot  of  the  second  column  can  be  melted  easily  with 
the  Bunsen  flame. 

The  methods  of  manufacture  and  the  treatment  of  metals  are 
much  influenced  also  by  their  volatility.  The  following  are  easily 
distilled:  Mercury,  b.-p.  357°;  potassium  and  sodium,  b.-p.  about 
700°;  cadmium,  b.-p.  770°;  zinc,  b.-p.  920°.  Even  the  most  in- 
volatile  metals  can  be  converted  into  vapor  in  the  electric  arc. 

In  many  cases  molten  metals  dissolve  in  one  another,  forming 
alloys.  Some  alloys  are  simply  solid  solutions.  Sometimes,  as 
in  the  case  of  lead  and  tin,  mixtures  can  be  formed  in  all  pro- 
portions. Again,  the  solubility  may  be  limited,  as  in  the  case  of 
zinc  and  lead,  where  only  1.6  parts  of  the  former  dissolve  in  100 
parts  of  the  latter.  Frequently  chemical  compounds  are  formed. 
The  colors  of  alloys  are  not  the  average  of  those  of  the  constitu- 
ents. Thus,  the  nickel  alloy  used  in  coining  contains  75  per  cent 
of  copper  and  25  per  cent  of  nickel,  yet  it  shows  none  of  the  color 
of  the  former. 

Alloys  in  which  mercury  is  one  of  the  components  are  known 
as  amalgams  (Gk.  /xaXay/wt,  a  soft  mass),  and  are  formed  with 
especial  ease  by  the  lighter  metals.  Of  the  common  metals,  iron 
is  the  least  miscible  with  mercury. 

The  good  conductivity  of  metals  for  electricity  distinguishes 
them  with  some  degree  of  sharpness  from  the  non-metals.  They 
show  considerable  variation  amongst  themselves,  silver  conducting 
sixty  times  as  well  as  mercury.  The  following  table  gives  the 
conductivities  of  the  metals,  expressed  in  terms  of  the  number  of 


436  COLLEGE    CHEMISTRY 

meters  of  wire  1  sq.  mm.  in  section  which,  at  15°,  offer  a  resistance 

of  one  ohm: 

Silver,  cast 62,89  Nickel,  cast 7.59 

Copper,  commercial  .    .  57.40  Iron,  drawn 7.55 

Gold,  cast 46.30  Platinum     ...'...  5.7-8.4 

Aluminium,  commercial  31.52  Steel     .    .    .    .V    .    .    .5.43 

Zinc,  rolled 16.95  Lead 4.56 

Brass 14.17  Mercury 1.049 

To  compare  these  conductivities  with  those  of  solutions,  it  may 
be  said  that  decinormal  hydrochloric  acid  (p.  240)  has  a  con- 
ductivity on  the  above  scale  of  0.035,  or  a  thirtieth  of  that  of 
mercury. 

The  world's  production  (1913)  of  the  metals  in  metric  tons  of 
1000  kilos,  is  approximately^  follows: 

Copper         1,000,000  Chromium  50,000  Gold          680 

Zinc  1,000,000  Nickel          32,000  Bismuth'  500 

Lead  1,000,000  Silver  7,800  Cadmium    50 

Tin  120,000  Tungsten       4,800  Platinum      9 

Aluminium       79,000  Mercury        3,000 

General  Chemical  Relations  of  the  Metallic  Elements.— 

Since  most  of  the  compounds  of  the  metals  are  ionogens,  their 
solutions,  except  when  the  metal  is  a  part  of  a  compound  ion,  all 
contain  the  metal  in  the  ionic  state,  and  the  resulting  substances, 
such  as  potassium-ion  and  cupric-ion,  have  constant  properties, 
irrespective  of  the  nature  of  the  negative  ion  with  which  they  may 
be  mixed.  The  properties  of  the  ions,  simple  and  compound,  are 
much  used  in  making  tests  in  analytical  chemistry.  On  the  other 
hand,  the  chemical  properties  of  the  oxides  and  of  the  salts  in  the 
dry  stale  are  of  importance  in  connection  with  metallurgy. 

There  are  three  chemical  properties  which  are  characteristic  of 
the  metallic  elements.  The  first  two  of  them  have  already  been 
discussed  somewhat  fully. 

1.  The  metals  are  able  by  themselves  to  form  positive  radicals 
of  salts,  and  therefore  to  exist  alone  as  positive  ions  (pp.  246,  296). 

2.  The  oxides  and  hydroxides  of  the  metals  are  basic   (pp. 
94,  296). 

3.  Each  typical  metal  has  at  least  one  halogen  compound  which 
is  little,  if  at  all,  hydrolyzed  by  water  (p.  296).     The  same  thing 
is  true  of  nitrates  and  other  salts  involving  active  acids. 


THE   BASE-FORMING   ELEMENTS  437 

In  refe'rence  to  the  third  characteristic,  the  non-hydrolysis  of 
halides  of  typical  metals,  a  word  of  explanation  is  required.  Active 
bases  (hydroxides  of  typical  metals),  such  as  sodium  hydroxide,  give, 
with  feeble  acids,  such  as  H2S  (p.  271),  H3P04  (p.  370),  H2C03 
(p.  383),  H2Si03  (p.  428),  and  H3B03  (p.  432),  salts  whose  solutions 
are  alkaline  in  reaction.  This  is  due  to  hydrolysis.  But  active 
bases  give,  with  active  acids,  such  as  HC1  and  HNO3,  salts  whose 
solutions  are  neutral  in  reaction.  This  is  the  fact  expressed  in  the 
third  characteristic  of  the  metallic  elements.  The  less  active  bases, 
being  hydroxides  of  less  active  metallic  elements,  give,  with 
active  acids,  salts  whose  solutions  are  not  neutral,  but  acid  in 
reaction.  Thus  cupric  chloride  solution  is  feebly  acid.  This  is 
because  there  is  a  tendency  for  the  ions  of  the  water  to  form  the 
slightly  dissociated  molecules  of  the  base : 

Cu++  -f  20H~  +  2H+  -»  Cu(OH)2  +  2H+. 

Finally,  a  salt  derived  from  a  base  and  an  acid,  both  of  which  are 
weak,  is  also  hydrolyzed.  If  the  resulting  base  or  acid  is  insoluble, 
the  hydrolysis  may  go  to  completion.  Aluminium  carbonate  and 
ammonium  silicate  (p.  429)  are  examples  of  salts  which,  for  this 
reason,  are  completely  hydrolyzed.  The  resulting  mixture  may 
have  an  acid  or  a  basic  reaction,  if  the  acid  or  the  base  is  sufficiently 
soluble  and  sufficiently  active.  Thus,  ammonium  sulphide  (NH^S 
solution  is  alkaline. 

Aside  from  these  points,  many  features  in  the  behavior  of 
metals  and  their  compounds  are  summed  up  in  the  electromotive 
series  (p.  260).  Before  proceeding  farther,  the  reader  should  re- 
read all  the  pages  referred  to  above.  He  should  also  reexamine 
the  various  kinds  of  chemical  changes  discussed  on  pp.  166,  197, 
251  et  seq.  and  particularly  the  varieties  of  ionic  chemical  change 
on  p.  259. 

Occurrence  of  the  Metals  in  Nature.  —  The  minerals  from 
which  metals  are  extracted  are  known  as  ores.  They  present  a 
comparatively  small  number  of  different  kinds  of  compounds. 
Most  of  the  metals  are  found  in  more  than  one  of  these  forms,  so 
that  in  the  following  statement  the  same  metal  frequently  occurs 
more  than  once. 

When  the  metal  occurs  free  in  nature,  it  is  said  to  be  native. 


438  COLLEGE    CHEMISTRY 

Thus  we  have  gold,  silver,  metals  of  the  platinum  group,  copper, 
mercury,  bismuth,  antimony,  and  arsenic  occurring  native  (cf. 
p.  60). 

The  metals  whose  oxides  are  important  minerals  are  iron,  man- 
ganese, tin,  zinc,  copper,  and  aluminium.  The  metals  are  ob- 
tained commercially  from  the  oxides  in  each  of  these  cases. 

The  metals  whose  sulphides  are  used  as  ores  are  iron,  nickel, 
cobalt,  antimony,  lead,  cadmium,  zinc,  copper,  and  mercury. 

From  the  carbonates  we  obtain  iron,  lead,  zinc,  and  copper. 
Several  other  metals,  such  as  manganese,  magnesium,  barium, 
strontium,  and  calcium  occur  in  larger  or  smaller  quantities  in  the 
same  form  of  combination. 

The  metals  which  occur  as  sulphates  are  those  whose  sulphates 
are  not  freely  soluble,  namely,  lead,  barium,  strontium,  and 
calcium. 

Compounds  of  metals  with  the  halogens  are  not  so  numerous. 
Silver  chloride  furnishes  a  limited  amount  of  silver.  Sodium  and 
potassium  chlorides  are  found  in  the  salt-beds. 

The  natural  silicates  are  very  numerous,  but  few  are  used 
for  the  preparation  of  the  metals.  Many  are  employed  for  other 
commercial  purposes,  kaolin  (p.  430)  being  a  conspicuous  example. 

Methods  of  Extraction  from  the  Ores. —  The  art  of  extract- 
ing metals  from  their  ores  is  called  metallurgy.  Where  the  metal 
is  native,  the  process  is  simple,  since  melting  away  from  the  matrix 
(p.  264)  is  all  that  is  required.  Frequently  a  flux  is  added.  A  flux 
usually  is  a  substance  which  interacts  with  infusible  materials 
to  give  fusible  ones.  It  combines  with  the  matrix,  giving  a  fusible 
slag  (resembling  glass).  Since  the  slag  is  a  melted  salt,  usually 
a  silicate,  and  does  not  mix  at  all  with  the  molten  metal,  separa- 
tion of  the  products  is  easily  effected.  When  the  ore  is  a  com- 
pound, the  metal  has  to  be  liberated  by  our  furnishing  a  material 
capable  of  combining  with  the  other  constituent.  The  details  of 
the  process  depend  on  various  circumstances.  Thus  the  volatile 
metals,  like  zinc  and  mercury,  are  driven  off  in  the  form  of  vapor, 
and  secured  by  condensation.  The  involatile  metals,  like  copper 
and  iron,  run  to  the  bottom  of  the  furnace  and  are  tapped  off. 

Where  the  ore  is  an  oxide  it  is  usually  reduced  by  heating  with 
carbon  in  some  form.  This  holds  for  the  oxides  of  iron  and  cop- 


THE   BASE-FORMING   ELEMENTS  439 

per.  Some  oxides  are  not  reducible  by  carbon  in  an  ordinary 
furnace.  Such  are  the  oxides  of  calcium,  strontium,  barium,  mag- 
nesium, aluminium,  and  the  members  of  the  chromium  group. 
At  the  temperature  of  the  electric  furnace  even  these  may  be  re- 
duced, but  the  carbides  are  formed  under  such  circumstances,  and 
the  metals  are  more  easily  obtained  otherwise.  Recently,  heat- 
ing the  pulverized  oxide  with  finely  powdered  aluminium  has  come 
into  use,  particularly  for  operations  on  a  small  scale.  Iron  oxide 
is  easily  reduced  by  this  means,  and  even  the  metals  manganese 
and  chromium  may  be  liberated  from  their  oxides  quite  readily 
by  this  action.  On  account  of  the  great  amounts  of  heat  liberated, 
this  procedure  has  received  the  name  aluminothermy  (q.v.). 

When  the  ore  is  a  carbonate,  it  is  first  heated  strongly  to  drive  out 
the  carbon  dioxide  (cf.  p.  381) :  FeCO3  ?±  FeO  +  C02t ,  and  then 
the  oxide  is  treated  according  to  one  of  the  above  mentioned 
methods.  When  the  ore  is  a  sulphide,  it  has  to  be  calcined  (p. 
275),  in  order  to  remove  the  sulphur,  and  the  resulting  oxide  is 
then  reduced. 

Chlorides  and  fluorides  of  the  metals  can  be  decomposed  by 
heating  with  metallic  sodium.  This  method  was  formerly  em- 
ployed in  the  making  of  magnesium  and  aluminium. 

The  metals  which  are  not  readily  secured  in  any  of  the  above 
ways  can  be  obtained  easily  by  electrolysis  of  the  fused  chloride 
or  of  some  other  compound.  Aluminium  is  now  manufactured 
entirely  by  the  electrolysis  of  a  solution  of  aluminium  oxide  in 
molten  cryolite. 

Compounds  of  the  Metals:    Oxides  and  Hydroxides. — 

The  oxides  may  be  made  by  direct  burning  of  the  metal,  by  heating 
the  nitrates  (cf.  p.  351),  the  carbonates  (cf.  p.  381),  or  the  hydrox- 
ides: Ca(OH)2<=>CaO  + H20t.  They  are  practically  insoluble 
in  water,  although  the  oxides  of  the  metals  of  the  alkalies  and  of 
the  metals  of  the  alkaline  earths  interact  with  water  rapidly  to 
give  the  hydroxides.  Oxides  are  usually  stable.  Those  of  gold, 
platinum,  mercury,  and  silver  decompose  when  heated,  yet  with 
increasing  difficulty  in  this  order.  The  metals,  like  the  non- 
metals,  frequently  give  several  different  oxides.  Those  of  the 
univalent  metals  (having  the  form  K20),  if'  we  leave  cuprous 
oxide  and  aurous  oxide  out  of  account,  have  the  most  strongly 


440  COLLEGE    CHEMISTRY 

basic  qualities.  Those  of  the  bivalent  metals  of  the  form  MgO, 
when  this  is  the  only  oxide  which  they  furnish,  are  base-forming. 
Those  of  the  trivalent  metals  of  the  form  A12O3,  known  as  ses- 
quioxides  (Lat.  sesqui-,  one-half  more),  are  the  least  basic  of  the 
basic  oxides.  The  oxides  of  the  forms  SnO2,  SbsA,  CrO3,  and 
Mn2O7,  in  which  the  metals  have  valences  from  4  to  7,  are  mainly 
acid-forming  oxides,  although  the  same  elements  usually  have 
other  lower  oxides,  which  are  basic. 

The  hydroxides  are  formed,  in  the  cases  of  the  metals  of  the 
alkalies  and  alkaline  earths,  by  direct  union  of  water  with  the 
oxides.  They  are  produced  also  by  double  decomposition  when 
a  soluble  hydroxide  acts  upon  a  salt  (cf.  p.  252).  All  hydroxides, 
except  those  of  the  alkali  metals,  lose  the  elements  of  water  when 
heated,  and  the  oxide  remains.  In  some  cases  the  loss  takes  place 
by  stages,  just  as  was  the  case  with  orthophosphoric  acid  (p.  368). 
Thus  lead  hydroxide  Pb(OH)2  (q.v.)  first  gives  the  hydroxide 
Pb2O(OH)2,  then  Pb302(OH)2,  and  then  the  oxide  PbO.  The 
hydroxides  of  mercury  and  silver,  if  they  are  formed  at  all,  are 
evidently  unstable,  for,  when  either  material  is  dried,  it  is  found 
to  contain  nothing  but  the  corresponding  oxide.  The  hydroxides, 
with  the  exception  of  those  of  the  metals  of  the  alkalies  and 
alkaline  earths,  are  all  little  soluble  in  water. 

Compounds  of  the  Metals:  Salts. —  It  may  be  said,  in 
general,  that  each  metal  may  form  a  salt  by  combination  with 
each  one  of  the  acid  radicals.  In  the  succeeding  chapters  we  shall 
describe  only  those  salts  which  are  manufactured  commercially,  or 
are  of  special  interest  for  some  other  reason.  The  various  salts 
will  be  described  under  each  metal.  Here,  however,  a  few  re- 
marks may  be  made  about  the  characteristics  of  the  more  com- 
mon groups  of  salts. 

The  chlorides  may  be  made  by  the  direct  union  of  chlorine  with 
the  metal  (cf.  p.  160),  or  by  the  combined  action  of  carbon  and 
chlorine  upon  the  oxide  (cf.  p.  426).  The  latter  method  is  used 
in  making  chromium  chloride.  The  general  methods  for  making 
any  salt  (p.  146),  such  as  the  interaction  of  a  metal  with  an  acid, 
or  of  the  oxide,  hydroxide,  or  another  salt  with  an  acid,  or  the 
double  decomposition  of  two  salts,  may  be  used  also  for  making 
chlorides.  The  chlorides  are  for  the  most  part  soluble  in  water. 


THE   BASE-FORMING   ELEMENTS  441 

Silver  chloride,  mercurous  chloride,  and  cuprous  chloride  are  al- 
most insoluble,  however,  and  lead  chloride  is  not  very  soluble. 
Most  of  the  chlorides  of  metals  dissolve  without  decomposition, 
but  hydrolysis  is  conspicuous  in  the  case  of  the  chlorides  of  the 
trivalent  metals,  such  as  aluminium  chloride  and  ferric  chloride 
(cf.  p.  437).  The  chlorides  of  some  of  the  bivalent  metals  are 
hydrolyzed  also,  but,  as  a  rule,  only  when  they  are  heated  with 
water.  This  is  the  case  with  the  chlorides  of  magnesium,  calcium, 
and  zinc.  Most  of  the  chlorides  are  stable  when  heated,  but  those 
of  the  noble  metals,  particularly  gold  and  platinum,  are  de- 
composed, and  chlorine  escapes.  The  chlorides  are  usually  the 
most  volatile  of  the  salts  of  a  given  metal,  and  so  are  preferred  for 
the  production  of  the  spectrum  of  the  metal.  Some  metals,  for  ex- 
ample iron,  form  two  or  more  different  chlorides.  Indium  gives 
InCl,  InCl2,  and  InCl3. 

The  sulphides  are  formed  by  the  direct  union  of  the  metal  with 
sulphur,  or  by  the  action  of  hydrogen  sulphide  or  of  some  soluble 
sulphide  upon  a  solution  of  a  salt  (cf.  p.  273).  In  one  or  two  cases 
they  are  made  by  the  reduction  of  the  sulphate  with  carbon.  The 
sulphides,  except  those  of  the  alkali  metals,  are  but  little  soluble 
in  water.  The  sulphides  of  aluminium  and  chromium  are  hy- 
drolyzed completely  by  water,  giving  the  hydroxides,  and  those 
of  the  metals  of  the  alkaline  earths  are  partially  hydrolyzed  (cf. 
p.  273). 

The  carbides  are  usually  formed  in  the  electric  furnace  by  inter- 
action of  an  oxide  with  carbon  (cf.  p.  379).  Some  of  them  are 
decomposed  by  contact  with  water,  after  the  manner  of  calcium 
carbide,  giving  a  hydroxide  and  a  hydrocarbon.  Of  this  class  are 
lithium  carbide  Li2C2,  barium  and  strontium  carbides  BaC2  and 
SrC2,  aluminium  carbide  AUCa,  manganese  carbide  MnC,  and  the 
carbides  of  potassium  and  glucinum.  Others,  such  as  those  of 
molybdenum  Mo2C  and  chromium  Cr3C2,  are  not  affected  by 
water. 

The  nitrates  may  be  made  by  any  of  the  methods  used  for 
preparing  salts.  They  are  all  at  least  fairly  soluble  in  water. 

The  sulphates  are  made  by  the  methods  used  for  making  salts, 
and  in  some  cases  by  the  oxidation  of  sulphides.  They  are  all 
soluble  in  water,  with  the  exception  of  those  of  lead,  barium,  and 
strontium.  Calcium  sulphate  is  meagerly  soluble. 


442  COLLEGE   CHEMISTRY 

The  carbonates  are  prepared  by  the  methods  used  for  making 
salts.  They  are  all  insoluble  in  water,  with  the  exception  of  those 
of  sodium  and  potassium.  The  hydroxides  of  aluminium  and  tin 
are  so  feebly  basic  that  they  do  not  form  stable  carbonates  (cf. 
pp.  429,  437). 

The  phosphates  and  silicates  are  prepared  by  the  methods  used 
in  making  salts.  The  former  are  obtained  also  by  special  processes 
already  described  (p.  371).  With  the  exception  of  the  salts  of 
sodium  and  potassium,  all  the  salts  of  both  these  classes  are 
insoluble. 

For  the  exact  solubilities  of  a  large  number  of  bases  and  salts 
at  18°,  see  the  Table  inside  the  cover,  at  the  front  of  this  book. 
Solubilities  at  all  temperatures  are  shown  in  the  diagram,  Fig.  58, 
p. 131. 

Exercises.  —  1.  What  do  we  mean  by  saying  that  an  oxide  is 
strongly  or  feebly  basic,  or  that  it  is  acidic? 

2.  What  is  meant  by  the  same  terms  when  applied  to  an 
hydroxide? 

3.  Compare  the  molar  solubilities  at  18°,  (a)  of  the  halides  of 
silver,  and  (6)  of  the  carbonates  and  (c)  oxalates  of  the  metals  of 
the  alkaline  earths,  noting  the  relation  between  solubility  and 
atomic  weight. 

4.  What  is  the  molar  concentration  of  chloride-ion  in  saturated 
solutions  of  silver  chloride  and  lead  chloride  at  18°,  assuming  com- 
plete ionization  in  these  very  dilute  solutions? 


CHAPTER  XXXIV 

THE   METALLIC   ELEMENTS   OF   THE  ALKALIES: 
POTASSIUM  AND  AMMONIUM 

THE  metals  of  this  family,  with  their  atomic  weights,  are: 

Lithium,. Li 6.9      Rubidium,  Rb    ......     85,5 

Sodium,  Na  (Ger.  natrium}      .    23.0      Caesium,  Cs         132.8 

Potassium,  K  (Ger.  kalium)      .    39.1 

The  Chemical  Relations  of  the  Metallic  Elements  of  the 
Alkalies.  —  The  metals  which  are  chemically  most  active  are 
included  in  this  group,  and  the  activity  increases  with  rising  atomic 
weight,  csesium  being  the  most  active  positive  element  of  all.  A 
freshly  cut  surface  of  any  of  these  metals  tarnishes  by  oxidation  as 
soon  as  it  is  exposed  to  the  air.  All  of  these  metals  decompose 
water  violently  (cf.  p.  60),  liberating  hydrogen.  The  hydroxides 
which  are  formed  by  this  action  are  exceedingly  active  bases,  that 
is  to  say,  they  give  a  relatively  large  concentration  of  hydroxide-ion 
in  solutions  of  a  given  molecular  concentration  (p.  243).  In  the 
dry  form  these  hydroxides  are  not  decomposed  by  heating,  while 
the  hydroxides  of  all  other  metals  lose  water  more  or  less  easily. 
In  all  their  compounds  the  metals  of  the  alkalies  are  univalent. 

The  compounds  of  ammonium  are  discussed  in  connection  with 
those  of  potassium,  to  which  they  present  the  greatest  resemblance. 

The  solubilities  are  often  decisive  factors  in  connection  with  the 
preparation  and  use  of  salts.  The  reader  will  find  most  of  these  in 
the  table  on  the  inside  of  the  cover,  at  the  front  of  this  book,  or  in 
the  diagram  on  p.  131,  and,  as  a  rule,  the  values  will  not  be  re- 
peated in  the  descriptive  paragraphs. 

POTASSIUM  K 

Occurrence.  —  Silicates  containing  potassium,  such  as  felspar 
and  mica  (p.  430),  are  constant  constituents  of  volcanic  rocks. 
These  minerals  are  not  used  commercially  as  sources  of  potassium 

443 


/[/[/[  COLLEGE    CHEMISTRY 

compounds.  The  salt  deposits  (see  below)  contain  potassium 
chloride,  alone  (sylvite)  and  in  combination  with  other  salts,  and 
most  of  the  compounds  of  potassium  are  manufactured  from  this 
material.  Part  of  our  potassium  nitrate,  however,  is  purified 
Bengal  saltpeter  (p.  347).  Potassium  sulphate  occurs  also  in  the 
salt  layers. 

Preparation.  —  Potassium  was  first  made  by  Davy  (1807)  by 
bringing  the  wires  from  a  battery  in  contact  with  a  piece  of  moist 
potassium  hydroxide.  Globules  of  the  metal  appeared  at  the 
negative  wire.  Electrolytic  processes  have  now  come  back  into 
use,  commercially,  molten  potassium  chloride  being  the  substance 
decomposed.  Castner's  reduction  process  involves  the  heating  of 
potassium  hydroxide  with  a  spongy  mass  of  carbide  of  iron  (CFe2). 
The  potassium  passes  off  as  vapor,  and  is  condensed: 
6KOH  +  2C  ->  2K2C03  +  3H2  +  2K. 

Physical  and  Chemical  Properties.  —  Potassium  is  a  silver- 
white  metal  (m.-p.  62°).  It  boils  at  720°,  giving  a  green  vapor. 

The  density  of  the  vapor  shows  the  molecular  weight  of  potas- 
sium to  be  about  40,  so  that  the  vapor  is  a  monatomic  gas.  The 
element  unites  violently  with  the  halogens,  sulphur,  and  oxygen. 
In  consequence  of  the  latter  fact  it  is  usually  kept  under  petroleum, 
an  oil  which  neither  contains  oxygen  itself,  nor  dissolves  a  sufficient 
amount  of  moisture  from  the  air  to  permit  much  oxidation  of  the 
potassium  to  take  place.  A  white,  crystalline  hydride  KH  is 
formed  when  hydrogen  is  passed  over  potassium  heated  to  360°. 
When  thrown  into  water  it  gives  potassium  hydroxide,  and  the 
hydrogen  is  liberated. 

Potassium  Chloride  KCl.  —  Sea-water  and  the  waters  of  salt 
lakes  contain  a  relatively  small  proportion  of  potassium  com- 
pounds. During  the  evaporation  of  such  waters,  however,  the 
potassium  compounds  tend  to  accumulate  in  the  mother-liquor 
while  sodium  chloride  is  being  deposited  on  the  bottom.  Hence 
the  upper  layers  of  salt  deposits  are  the  richest  in  compounds  of 
potassium.  Thus,  at  Stassfurt,  near  Magdeburg,  there  is  a  thick- 
ness of  more  than  a  thousand  meters  of  common  salt.  Above  this 
are  25-30  meters  of  salt  layers  in  which  the  potassium  salts  are 
chiefly  found. 


POTASSIUM  445 

The  chief  forms  in  which  potassium  chloride  is  found  in  the  salt 
beds  are  sylvite  KC1  and  carnallite  KCl,MgCl2,6H2O.  The  latter 
salt  is  heated  with  a  small  amount  of  water,  or  with  a  mother- 
liquor  obtained  from  a  previous  operation  and  containing  sodium 
and  magnesium  chlorides.  From  the  clear  liquid,  when  it  cools, 
potassium  chloride  is  deposited  first  and  then  carnallite.  The 
former  is  taken  out  and  purified,  and  the  latter  goes  through  the 
process  again.  This  potassium  chloride  is  the  source  from  which 
our  other  potassium  compounds  are  made.  It  is  also  our  chief 
potassium-bearing  fertilizer.  It  is  a  white  substance  crystallizing 
in  cubes,  melting  at  about  750°,  and  slightly  volatile  at  high  tem- 
peratures. 

Recently,  the  giant  kelps  of  the  Pacific  coast  have  been  used  as 
a  source  of  potassium  chloride.  The  dried  seaweed  contains  9  per 
cent  of  this  salt  and  about  0.1  per  cent  of  iodine. 

The  Other  Halides  of  Potassium.  —  When  iodine  is  heated 
in  a  strong  solution  of  potassium  hydroxide,  potassium  iodate  and 
potassium  iodide  are  both  formed  (p.  318): 

6KOH  +  3I2  -»  5KI  +  KI03  +  3H20. 

The  dry  residue  from  evaporation  is  heated  with  powdered  carbon 
to  reduce  the  iodate,  and  all  the  iodide  can  then  be  purified  by 
recrystallization.  The  salt  forms  large,  somewhat  opaque  cubes 
(m.-p.  623°).  It  is  used  in  medicine  and  for  precipitating  silver 
iodide  Agl  in  photography  (q.v.). 

The  aqueous  solution  takes  up  free  iodine,  forming  KI3,  in 
equilibrium  with  dissolved  iodine:  Is~^  I"  +  ^2  (dslvd).  It  is 
used  in  testing  for  starch,  and  in  reactions  in  which  a  solution  of  free 
iodine  is  required. 

Potassium  bromide  KBr  may  be  made  in  the  same  way  as 
the  iodide.  It  crystallizes  in  cubes.  It  is  used  in  medicine  and 
for  precipitating  silver  bromide  in  making  photographic  plates 
(q*.). 

The  fluoride  of  potassium  K2F2  may  be  obtained  by  treating  the 
carbonate  or  hydroxide  with  hydrofluoric  acid.  It  is  a  deliques- 
cent, white  salt.  When  treated  with  an  equimolecular  quantity 
of  hydrofluoric  acid  it  forms  potassium-hydrogen  fluoride  KHF2,  a 
white  salt  which  is  also  very  soluble. 


446 


COLLEGE    CHEMISTRY 


Potassium  Hydroxide  KOH.  —  This  compound,  known  also 
as  caustic  potash,  and  colloquially  as  potassium  hydrate,  was 
formerly  made  entirely  by  boiling  potassium  carbonate  with  cal- 
cium hydroxide  suspended  in  water  (milk  of  lime)  : 


The  operation  is  conducted  in  iron  vessels,  because  porcelain,  being 
composed  of  silicates,  interacts  with  solutions  of  bases.  On 
account  of  the  very  limited  solubility  of  the  calcium  hydroxide 
(0.17  g.  in  100  g.  Aq),  the  water  takes  up  fresh  portions  into  solution 
only  when  the  part  dissolved  has  already  undergone  chemical 
change.  The  calcium  carbonate  which  is  precipitated  is,  however, 
still  more  insoluble  (0.0013  g.  in  100  g.  Aq),  and  hence  the  action 
goes  forward.  After  the  precipitate  has  settled,  the  potassium 
hydroxide  is  obtained  by  evaporation  of  the  clear  liquid,  K+  + 
GET  -»  KOH. 

Potassium    hydroxide    is    now    manufactured    by    electrolytic 
processes.     When  a  solution  of  potassium  chloride  is  electrolyzed, 


FIG.  112. 

chlorine  is  liberated  at  the  anode,  and  hydrogen  and  potassium 
hydroxide  at  the  cathode  (p.  228).  These  two  sets  of  products 
must  be  kept  apart,  since  by  their  interaction  potassium  hypo- 
chlorite  and  potassium  chloride  would  be  formed  (cf.  p.  308).  In 
the  Castner-Kellner  apparatus  (Fig.  112),  which  serves  for  making 
either  potassium  or  sodium  hydroxide,  the  two  end  compartments 
are  filled  with  potassium  chloride  solution  (or  brine)  and  contain 
the  graphite  anodes.  The  central  compartment  contains  potas- 
sium hydroxide  solution  and  the  iron  cathode.  The  positive 


POTASSIUM  447 

current  enters  by  the  anodes,  and  the  chlorine  is  therefore  at- 
tracted to  and  liberated  upon  the  graphite:  2C1~  +  2©  — >  C12. 
After  rising  through  the  liquid,  it  is  collected  for  the  manufacture 
of  liquefied  chlorine  or  of  bleaching  powder.  The  ions  of  potassium 
(or  of  sodium)  are  discharged  upon  a  layer  of  mercury  which  covers 
the  whole  floor  of  the  box,  and  the  free  metal  dissolves  in  the  mer- 
cury, forming  an  amalgam  (p.  435).  The  layer  of  mercury  extends 
beneath  the  partitions,  and  a  slight  rocking  motion  given  to  the 
cell  causes  the  amalgam  to  flow  below  the  partition  into  the  central 
compartment.  Here  the  potassium  leaves  the  mercury  in  the 
form  of  potassium-ion  and  is  attracted  by  the  cathode.  Upon 
this,  hydrogen  from  the  water  is  discharged,  and  the  residual 
hydroxide-ion,  together  with  the  metal-ion,  constitute  potassium 
or  sodium  hydroxide : 

2K+  +  2H+  +  20H~  +  20  ->  2K+  +  20H~  +  H2. 

A  slow  influx  of  salt  solution  to  the  end  compartments,  and  over- 
flow of  the  alkaline  solution  in  the  central  cell,  are  maintained. 
The  overflowing  liquid  contains  20  per  cent  of  the  alkali.  Since 
there  is  no  undecomposed  chloride  present  in  the  part  of  the  solu- 
tion which  contains  the  hydroxide,  simple  evaporation  to  dryness 
furnishes  the  solid  alkali.  Other  forms  of  electrolytic  cells,  such  as 
the  Briggs,  and  the  Townsend-Baekeland  cells,  are  also  largely 
in  use. 

Potassium  hydroxide  is  exceedingly  soluble  in  water,  and  conse- 
quently, instead  of  being  crystallized  from  solution,  the  molten 
residue  from  evaporation  is  cast  in  sticks.  The  hydroxide  is  highly 
deliquescent.  It  also  absorbs  carbon  dioxide  from  the  air,  giving 
potassium  carbonate.  Solutions  of  the  hydroxide  have  an  exceed- 
ingly corrosive  action  upon  the  flesh,  resolving  it  into  a  slimy 
mass  by  decomposing  the  proteins.  In  solution,  the  base  is  highly 
ionized,  furnishing  a  high  concentration  of  hydroxide-ion.  Com- 
mercially, it  is  chiefly  employed  in  the  making  of  soft  soap. 

Potassium  oxide  K2O  may  be  made  by  heating  potassium  nitrate 
with  potassium  in  a  vessel  from  which  air  is  excluded :  2KN03  -f 
10K  — >  6K2O  +  N2.  It  interacts  violently  with  water,  giving  the 
hydroxide.  When  exposed  to  the  air  it  unites  spontaneously  with 
oxygen,  and  a  yellow  peroxide  K204  is  formed.  The  same  peroxide 
is  formed  when  potassium  burns  in  air  or  oxygen. 


448  COLLEGE    CHEMISTRY 

Potassium  Chlorate  KCIO3.  —  The  preparation  of  this  salt 
by  interaction  of  potassium  chloride  with  calcium  chlorate  has 
already  been  described  (p.  313).  It  is  also  made  by  electrolysis  of 
potassium  chloride  solution,  the  potassium  hydroxide  and  chlorine 
which  are  liberated  being  precisely  the  materials  required.  All 
that  is  necessary  is  to  use  a  warm,  concentrated  solution  and  to 
provide  for  the  mixing  of  the  materials  generated  at  the  electrodes. 
The  salt  crystallizes  out  when  the  solution  cools. 

Potassium  chlorate  crystallizes  in  monoclinic  plates.  It  melts 
at  about  334°,  and  at  a  temperature  slightly  above  this  the  visible 
liberation  of  oxygen  begins  (cf.  pp.  27,  29).  On  account  of  the 
ease  with  which  its  oxygen  is  liberated,  the  salt  is  employed  in 
making  fireworks  and  as  a  component  of  the  heads  of  Swedish 
matches.  It  is  also  used  in  medicine. 

Potassium  perchlorate  KC1C>4,  formed  by  the  heating  of  the  chlo- 
rate (p.  315),  gives  white  crystals  belonging  to  the  rhombic  system. 

By  adding  chlorine-water  to  potassium  carbonate  solution,  a 
mixture  of  the  chloride  and  potassium  hypochlorite  is  formed : 

HC1  +  HC10  +  K2C03  4  KC1  +  KC1O  +  H2O  +  C02. 

The  carbonic  acid,  however,  is  not  completely  displaced  by  the 
HC10,  which  is  a  feeble  acid.  Hence,  the  solution  is  used,  under 
the  name  eau  de  Javel  (often  misspelt  Javelle),  in  the  household  for 
removing  stains. 

The  mode  of  preparing  potassium  bromate  KBr03  and  potassium 
iodate  KIOs  has  already  been  described  (p.  318).  Potassium 
iodate  may  be  made  also  very  conveniently  by  melting  together 
potassium  chlorate  and  potassium  iodide  at  a  low  temperature. 
The  iodate  is  much  less  soluble  (see  Table)  than  the  chloride,  and 
the  mixture  may  be  separated  by  crystallization  from  water. 

Potassium  Nitrate  KNO3.  —  The  formation  of  this  salt  in 
nature  and  its  mode  of  extraction  and  purification  have  already 
been  described  (p.  347).  This  source  of  supply  proved  insufficient, 
for  the  first  time,  during  the  Crimean  war  (1852-55),  and  a  method 
of  manufacture  from  Chile  saltpeter  (sodium  nitrate),  which  is  a 
much  cheaper  substance,  was  introduced.  Sodium  nitrate  and 
potassium  chloride  are  heated  with  very  little  water,  and  the  sodium 
chloride  produced  by  the  action,  which  is  a  reversible  one,  is  by 


POTASSIUM  449 

far  the  least  soluble  of  the  four  salts  (see  Diagram,  p.  131).  On 
the  other  hand,  in  the  hot  water,  the  potassium  nitrate  is  by  far 
the  most  soluble.  Hence  the  hot  liquid,  quickly  drained  from  the 
crystals  through  canvas,  contains  the  required  salt,  and  most  of 
the  sodium  chloride  is  in  the  form  of  a  precipitate.  If  the  solu- 
bility curve  of  potassium  nitrate  (p.  131)  is  examined,  it  will  be 
seen  that  this  salt  is  but  slightly  soluble  in  cold  water,  and  hence 
most  of  it  is  deposited  when  the  solution  cools. 
The  crystals  are  mixed  with  little  sodium  chloride, 
for,  as  the  curve  shows,  common  salt  is  little  less 
soluble  at  10°  than  it  is  at  100°. 

Potassium  nitrate  gives  long  prisms  belonging 
to  the  rhombic  system  (Fig.  113).  It  melts  at 
about  340°,  and  when  more  strongly  heated  gives 
off  oxygen,  leaving  potassium  nitrite  (p.  356). 
Although  it  does  not  form  a  hydrate,  the  crys- 
tals enclose  small  portions  of  the  mother-liquor, 
and  consequently  contain  both  water  and  im- 
purities. When  heated,  the  crystals  fly  to  pieces 
explosively  (decrepitate),  on  account  of  the  vapor- 
ization of  this  water.  Many  substances  which 
form  large  crystals  and  do  not  melt  at  a  low  temperature,  behave 
in  the  same  way  and  for  the  same  reason.  In  consequence  of 
this,  the  purest  salt  is  made  by  violent  stirring  of  the  solution 
during  the  operation  of  crystallization,  the  result  being  the  forma- 
tion of  a  crystal-meal. 

Potassium  nitrate  is  used  chiefly  in  the  manufacture  of  gun- 
powder, which  contains  75  per  cent  of  the  highly  purified  salt. 
The  other  components  are  10  per  cent  of  sulphur.,  14  per  cent  of 
charcoal,  and  about  1  per  cent  of  water.  The  ingredients  are 
intimately  mixed  in  the  form  of  paste,  and  the  material  when  dry 
is  broken  up  and  sifted,  grains  of  different  sizes  being  used  for 
different  purposes.  The  chemical  action  which  takes  place  when 
gunpowder  is  fired  in  an  open  space  gives  chiefly  potassium  sul- 
phide, carbon  dioxide,  and  nitrogen: 

2KNO3  +  3C  +  S  ->  K2S  +  3C02  +  N2. 

The  explosion  occurring  in  firearms  follows  a  much  more  complex 
course,  and  half  of  the  solid  product  is  said  to  be  potassium  car- 


450  COLLEGE    CHEMISTRY 

bonate  (a  solid,  hence  the  smoke).  One  gram  yields  264  c.c.  of 
gases  (0°  and  760  mm.),  and  a  much  larger  volume  at  the  tempera- 
ture of  the  explosion,  and  gives  660  calories.  The  pressure,  at  the 
temperature  of  the  explosion,  if  the  gases  could  be  confined  within 
the  volume  originally  occupied  by  the  gunpowder,  would  reach 
about  forty-four  tons  per  square  inch.  In  recent  years,  except  in 
mining,  common  gunpowder  has  been  displaced  largely  by  smoke- 
less powder  (pp.  358,  359),  which,  in  decomposing,  produces  no 
solids. 

Potassium  nitrate  is  used  also  in  preserving  ham  and  corned 
beef,  on  which  it  confers  a  red  color. 

Potassium  Carbonate  K2CO3.  —  This  salt  is  manufactured 
from  potassium  chloride,  from  the  Stassfurt  deposits.  The  chlo- 
ride is  heated  with  magnesium  carbonate  (magnesite),  water,  and 
carbon  dioxide  under  pressure: 

2KC1  +  3MgC03  +  C02  +  5H20  -» 2KHMg(C03)2,4H20  +  MgCl2. 

The  hydrated  mixed  salt  separates  from  the  liquid  containing 
magnesium  chloride  and  is  decomposed  by  heating  with  water  at 
120°.  The  product  is  a  solution  of  potassium  carbonate,  from 
which  the  precipitated  magnesium  carbonate  is  removed  by  filtra- 
tion and  used  over  again.  In  some  districts  potassium  carbonate 
is  still  extracted  from  wood-ashes,  its  original  source  and  the  origin 
of  its  name,  potash.  The  sugar  beet  takes  up  a  considerable 
amount  of  potash  from  the  soil,  and  the  extract,  after  removal  of 
the  sugar,  is  evaporated  and  calcined.  Wool  scourings,  when  evap- 
orated and  calcined,  also  afford  a  small  supply. 

This  salt  is  usually  sold  in  the  form  of  an  anhydrous  powder 
(m.-p.  over  1000°).  When  crystallized  from  water  it  gives  a 
hydrate  2K2CO3,3H20.  It  is  deliquescent.  Its  aqueous  solution, 
like  that  of  sodium  carbonate  (cf.  p.  383),  has  a  marked  alkaline 
reaction.  The  commercial  name  of  the  substance  is  pearl  ash. 
It  is  used  in  making  soft  soap  and  hard  (i.e.,  difficultly  fusible) 
glass.  It  is  also  employed,  by  interaction  with  acids,  in  making 
salts  of  potassium. 

The  use  of  the  bicarbonate  KHCO3  in  purifying  carbon  dioxide 
has  already  been  mentioned  (p.  381).  Before  the  nineteenth 
century,  this  salt  was  used  under  the  name  saleratus  (Lat.  aerated 


POTASSIUM  451 

salt),  a  name  now  sometimes  given  the  baking  soda  NaHC03 
which  has  displaced  it. 

Potassium  Cyanide  KNC.  —  This  salt  is  made  by  heating 

dry  potassium  ferrocyanide  (q.v.) : 

K4Fe(CN)6  -*  4KNC  +  Fe  +  2C  +  N2. 

When  the  residue  is  extracted  with  water,  only  the  potassium 
cyanide  dissolves,  and  it  is  easily  crystallized  in  pure  form  from  the 
solution. 

Potassium  cyanide  is  extremely  soluble  in  water,  and  is  therefore 
deliquescent.  Its  poisonous  qualities  are  equal  to  those  of  hydro- 
cyanic acid.  The  acid  is  so  feeble  as  to  be  liberated  both  by  the 
moisture  and  by  the  carbon  dioxide  of  the  air,  and  hence  this  salt 
always  has  an  odor  of  hydrocyanic  acid.  Potassium  cyanide  was 
used  in  electroplating  (q.v.),  and  in  extracting  gold  (q.v.)  from 
its  ores,  but  has  been  displaced  by  sodium  cyanide  NaNC,  which 
is  now  less  expensive. 

The  preparation  of  potassium  cyanate  KCNO,  a  white,  easily 
soluble  salt,  and  of  potassium  thiocyanate  KCNS,  a  white,  deli- 
quescent salt,  have  already  been  described  (p.  421). 

The  Sulphate  and  Bisulphate.  —  Potassium  sulphate  K2S04 
is  a  constituent  of  several  double  salts  found  in  the  Stassfurt  de- 
posits. It  is  extracted  from  schoenite  MgSO4,K2S04,6H20  and 
kainite  MgS04,MgCl2,K2SO4,6H2O.  The  former  is  treated  with 
potassium  chloride  and  comparatively  little  water,  whereupon  the 
relatively  insoluble  potassium  sulphate  crystallizes  out,  and  the 
magnesium  chloride  remains  in  the  mother-liquor.  The  crystals 
belong  to  the  rhombic  system,  contain  no  water  of  crystallization, 
and  melt  at  1066°.  This  salt  is  employed  in  preparing  alum  (q.v.) 
and  is  much  used  as  a  fertilizer.  Since  plants  take  up  solutions 
through  their  cell  walls,  they  can  absorb  soluble  compounds  only. 
They  are,  therefore,  dependent,  for  the  potassium  compounds 
which  they  require,  upon  the  weathering  out  of  soluble  potassium 
compounds  from  insoluble  silicates  containing  potassium  (p.  430) 
found  in  the  soil.  The  weathering  takes  place  too  slowly  to 
furnish  a  sufficient  supply  for  many  crops,  particularly  that  of  the 
sugar-beet.  Hence  potassium  sulphate  is  mixed  directly  with 
the  soil. 


452  COLLEGE    CHEMISTRY 

Potassium-hydrogen  sulphate  (bisulphate)  KHS04  is  made  by  the 
action  of  sulphuric  acid  upon  potassium  sulphate:  K2S04  +  H2SO4 
— >  2KHSO4.  It  crystallizes  from  water,  in  which  it  is  very  soluble, 
in  tabular  crystals.  Its  properties  are  similar  to  those  of  sodium 
bisulphate,  which  have  already  been  described  (p.  288). 

Sulphides  of  Potassium.  —  By  the  treatment  of  a  solution  of 
potassium  hydroxide  with  excess  of  hydrogen  sulphide,  a  solution  of 
potassium-hydrogen  sulphide  is  obtained.  Evaporation  of  the  solu- 
tion gives  a  deliquescent,  solid  hydrate  2KHS,H20.  When  the 
solution,  before  evaporation,  is  treated  with  an  equivalent  amount 
of  potassium  hydroxide,  and  the  water  is  driven  off,  potassium 
sulphide  K2S  remains  behind  (cf.  p.  270) : 

KHS  +  KOH  <±  K2S  +  H20. 

Considerable  amounts  of  sulphur  can  be  dissolved  in  solutions  of 
either  of  these  sulphides.  By  evaporation  of  the  resulting  yellow 
liquids,  various  polysulphides  have  been  obtained.  These  are 
probably  K2S5,  or  mixtures  of  the  pentasulphide  with  K2S  (cf.  p. 
274).  Similar  substances  are  produced,  as  a  result  of  the  libera- 
tion and  recombination  of  sulphur,  when  the  solutions  are  exposed 
to  the  oxidizing  action  of  the  air : 

2KHS  +  02  ->  2KOH  +  2S. 

Properties  of  Potassium-ion  K+:  Analytical  Reactions.  — 

The  positive  ionic  material  of  the  potassium  salts  is  a  colorless 
substance.  It  unites  with  all  negative  ions,  and  most  of  the 
resulting  compounds  are  fairly  soluble.  For  its  recognition  we  add 
solutions  containing  those  ions  which  give  with  it  the  least  soluble 
salts.  Thus,  with  chloroplatinic  acid  H2PtCl6  it  gives  a  yellow 
precipitate  of  potassium  chloroplatinate  K2PtCl6.  Since  nearly 
one  part  of  this  salt  dissolves  in  100  parts  of  water,  the  test  is  far 
from  being  a  delicate  one.  Picric  acid  (p.  349)  gives  potassium 
picrate  KC6H2(N02)3O,  which  is  much  less  soluble  in  water  (0.4 
parts  in  100  at  15°).  Perchloric  acid  and  hydrofluosilicic  acid 
likewise  give  somewhat  insoluble  salts  of  potassium.  Potassium- 
hydrogen  tartrate  KHC4H406  is  precipitated  by  the  addition  of 
tartaric  acid  to  a  sufficiently  concentrated  solution  of  a  potassium 


AMMONIUM  453 


salt.  The  neutral  tartrate  K^C-iHiOG  is  much  more  soluble.  The 
latter  may  be  obtained  by  treating  the  precipitate  with  a  solution 
of  potassium  hydroxide.  Addition  of  an  acid  to  this  solution 
causes  reprecipitation  of  the  bitartrate. 

A  much  more  delicate  test  for  the  recognition  of  a  potassium 
compound  consists  in  the  examination  by  means  of  the  spectroscope 
of  the  light  given  out  by  a  Bunsen  flame,  in  which  a  little  of  the 
salt  is  held  upon  a  platinum  wire.  When  the  amount  of  potassium 
is  considerable,  and  no  other  substance  which  would  likewise  color 
the  flame  is  present  to  mask  the  effect,  the  violet  tint  is  recognizable 
by  the  eye. 

Rubidium  and  Csesium.  —  In  1860  Bunsen  discovered 
several  new  lines  in  the  spectrum  given  by  materials  derived  from 
the  salts  in  Durkheim  mineral  water.  Two  new  elements  of  the 
alkali  group  were  found  to  cause  their  presence,  and  were  named, 
from  the  colors  of  the  lines  which  they  gave,  rubidium  (red)  and 
caesium  (blue).  Rubidium  is  obtainable  with  relative  ease  from 
the  mother-liquors  of  the  Stassfurt  works. 

The  metals  may  be  obtained  by  heating  their  hydroxides  with 
magnesium  powder.  The  hydroxides  of  these  two  elements  are 
more  active  as  bases  than  is  potassium  hydroxide.  Their  salts  are 
very  much  like  those  of  potassium. 

AMMONIUM 

The  compounds  of  ammonium  claim  a  place  with  those  of  the 
alkali  metals  because  in  aqueous  solution  they  give  ammonium-ion 
NH4+,  a  substance  which  in  its  behavior  closely  resembles  potas- 
sium-ion. Some  of  the  special  properties  peculiar  to  ammonium 
compounds,  and  particularly  the  properties  of  ammonium  hydrox- 
ide NKiOH,  have  been  discussed  in  detail  already  (pp.  343-345). 

Salts  of  Ammonium.  —  Ammonium  chloride  NH4C1,  known 
commercially  as  salammoniac,  like  all  the  other  compounds  of 
ammonium,  is  prepared  from  the  ammonia  dissolved  by  the  water 
used  to  wash  illuminating  gas  (p.  411),  or  that  obtained  from  by- 
product coke  ovens  (p.  411).  It  is  purified  by  sublimation,  and 
then  forms  a  compact  fibrous  mass.  At  337.8°  its  vapor  exercises 


454  COLLEGE    CHEMISTRY 

one  atmosphere  pressure,  and  is  dissociated  into  ammonia  and 
hydrogen  chloride  to  the  extent  of  62  per  cent  (p.  345). 

Ammonium  nitrate  NKiNOs  is  a  white  crystalline  salt  which 
may  be  made  by  the  interaction  of  ammonium  hydroxide  and 
nitric  acid.  When  heated  gently  (rn.-p.  166°)  it  decomposes,  giving 
nitrous  oxide  and  water  (p.  357).  It  is  used  as  an  ingredient  in 
fireworks  and  explosives. 

When  ammonium  hydroxide  is  treated  with  excess  of  carbon 
dioxide  the  solution  gives,  on  evaporation,  ammonium  bicarbonate 
NKiHCOa.  This  is  a  white  crystalline  salt  which  is  fairly  stable 
at  the  ordinary  temperature.  It  has,  however,  a  faint  odor  of 
ammonia,  and  its  dissociation  becomes  very  rapid  when  slight  heat 
is  applied.  When  a  solution  of  this  salt  is  treated  with  ammonium 
hydroxide,  the  normal  carbonate  (NH^COs  is  formed.  But  this 
salt,  when  left  in  an  open  vessel,  loses  ammonia  very  rapidly,  and 
leaves  the  bicarbonate  behind. 

Ammonium  thiocyanate  NHiNCS  (cf.  p.  421)  is  a  white  salt 
which  finds  some  application  in  analysis. 

Ammonium  sulphate  (NH^SC^  is  a  white  salt  which  is  used 
chiefly  as  a  fertilizer.  By  electrolysis  of  a  concentrated  solution 
of  the  bisulphate  NHiHSCX,  ammonium  persulphate  (NH4)2S2O8, 
which  is  less  soluble,  is  formed  and  crystallizes  out  (cf.  p.  291). 

Solutions  of  ammonium-hydrogen  sulphide  NH^HS  and  ammo- 
nium sulphide,  (NH^S,  made  by  passing  hydrogen  sulphide  gas  into 
ammonium  hydroxide,  are  much  used  in  analysis.  The  sulphide  is 
almost  completely  hydrolyzed  by  water  into  the  acid  sulphide  and 
ammonium  hydroxide,  its  behavior  being  like  that  of  sodium 
sulphide  (p.  271)  : 


2NH3  +  H2S  <=±  (NH4)2S  fc>  2NH4+  +  S=)  <_  WQ- 
H20±=>     OH-  +  H+J- 

It  is  used  for  the  precipitation  of  sulphides,  such  as  zinc  sulphide 
ZnS,  which  are  insoluble  in  water.  Although  the  S-  ions  are  not 
numerous  at  any  moment,  disturbance  of  the  equilibrium  by  their 
removal,  when  they  pass  into  combination,  causes  displacements 
which  result  in  the  generation  of  a  continuous  supply.  The  liquid 
smells  strongly  of  ammonia  and  hydrogen  sulphide,  on  account  of 
the  dissociation  of  the  parent  molecules  by  reversal  of  the  above 
equilibria. 


AMMONIUM  455 

The  solutions,  when  pure,  are  colorless.  They  dissolve  free 
sulphur,  giving  yellow  polysulphides  similar  to  those  of  potassium 
(p.  452).  The  same  yellow  substances  are  also  obtained  by 
gradual  oxidation  of  ammonium  sulphide,  when  the  solution  of  this 
salt  is  allowed  to  stand  in  a  bottle  from  which  the  air  is  imperfectly 
excluded. 

Ammonium  Amalgam.  —  When  a  salt  of  ammonium  is  de- 
composed by  electrolysis  the  NILf1",  upon  its  discharge,  ordinarily 
gives  ammonia  and  hydrogen,  and  no  substance  NHi  is  obtained. 
If,  however,  a  pool  of  mercury  is  used  as  the  negative  electrode,  the 
NH4  forms  an  amalgam  with  it,  and  there  seems  to  be  no  doubt  that 
this  substance  is  actually  present  in  solution  in  the  mercury. 
While  the  amalgam  is  being  formed  it  swells  up  and  gives  off  the 
decomposition  products  above  mentioned,  so  that  the  existence  of 
the  substance  is  only  temporary.  The  same  material  may  be 
obtained  by  putting  sodium  amalgam  into  a  strong  solution  of  a 
salt  of  ammonium.  The  action  is  a  displacement  of  one  ion  by 
another  (p.  259)  : 

Na(dslvd  in  mercury)  +  NH4+  —  >  NELi(dslvd  in  mercury)  +  Na+. 

This  behavior  is  interesting  since  it  is  in  harmony  with  the  idea  that 
ammonium,  if  it  could  be  isolated,  would  have  the  properties  of  a 
metal.  Substances  other  than  metals  are  not  miscible  with 
mercury. 


Ammonium-ion  NH^:  Analytical  Reactions.  —  Ionic  am- 
monium is  a  colorless  substance.  It  unites  with  negative  ions,  giv- 
ing salts,  which,  in  the  majority  of  cases,  are  soluble.  Ammonium 
chloroplatinate  (NH^PtCle,  and  to  a  less  extent  ammonium- 
hydrogen  tartrate  NttiH^KiOe,  are  insoluble  compounds,  and 
their  precipitation  is  used  as  a  test.  The  surest  means  of  recogniz- 
ing ammonium  compounds,  however,  consists  in  adding  a  soluble 
base  to  the  substance  (cf.  p.  345).  The  ammonium  hydroxide, 
which  is  thus  -formed,  gives  off  ammonia,  and  the  latter  may  be 
detected  by  its  odor. 

Exercises.  —  1.  What  kind  of  metals  will,  in  general,  interact 
with  solutions  of  bases  (cf.  p.  296)? 


456  COLLEGE   CHEMISTRY 

2.  Why  should  a  mixture  of  potassium  chlorate  and  antimony 
trisulphide  be  explosive? 

3.  How  should  you  set  about  making,  (a)  a  borate  of  potassium, 
(6)  potassium  pyrophosphate,  (c)  ammonium  nitrite,   (d)  ammo- 
nium chlorate,  (e)  ammonium  iodide? 


CHAPTER  XXXV 

SODIUM    AND    LITHIUM.     IONIC    EQUILIBRIUM    CONSIDERED 
QUANTITATIVELY 

SODIUM  chloride  forms  more  than  two-thirds  of  the  solid  matter 
dissolved  in  sea-water,  and  the  great  salt  deposits  are  largely  com- 
posed of  it.  Sea-plants  contain  mainly  sodium  salts  of  organic 
acids,  just  as  land-plants  contain  potassium  salts.  Chile  salt- 
peter and  albite  (a  soda  feldspar)  are  important  minerals. 

Compounds  of  sodium  are  usually  cheaper  than  the  correspond- 
ing ones  of  potassium.  Also,  since  the  atomic  weight  of  sodium  is 
23,  against  39  for  potassium,  a  smaller  weight  of  the  sodium  com- 
pound will  produce  the  same  chemical  result.  For  these  two 
reasons,  sodium  compounds,  except  in  special  cases, 
are  always  used  for  commercial  purposes. 

Preparation.  —  Sodium  was  first  made  by  Davy 
(1807)  by  electrolysis  of  moist  sodium  hydroxide. 
It  is  manufactured  by  the  electrolysis  of  fused 
sodium  hydroxide  by  a  method  invented  by  Castner. 
The  negative  electrode  projects  through  the  bottom 
of  the  iron  vessel  containing  the  fused  hydroxide 
(Fig.  114),  and  here  the  sodium  and  hydrogen  are 
liberated.  This  electrode  is  surrounded  by  a  wire-gauze  parti- 
tion to  permit  circulation  of  the  fused  mass,  but  prevent  escape 
of  the  globules  of  sodium.  This  is  surmounted  by  a  bell-shaped 
vessel  of  iron.  The  positive  electrode  is  an  iron  cylinder  sur- 
rounding the  gauze.  The  sodium  and  hydrogen  liberated  at  the 
cathode,  being  lighter  than  the  fused  mass,  ascend  into  the  iron 
vessel  (at  A),  under  the  edge  of  which  the  hydrogen  escapes. 
Oxygen  is  set  free  at  the  anode.  The  top  is  closed,  to  prevent 
the  sodium  from  burning.  The  melted  sodium  is  ladled  into 
molds,  like  candle  molds. 

457 


458  COLLEGE    CHEMISTRY 

Properties.  —  Sodium  is  a  soft,  shining  metal,  melting  at  96° 
and  boiling  at  742°.  The  green  vapor  is  a  monatomic  gas.  The 
general  chemical  properties  have  already  been  given  (p.  443). 
The  metal  unites  with  hydrogen  to  form  a  hydride  NaH,  which 
resembles  potassium  hydride  (p.  444).  The  amalgam  with  mer- 
cury, when  it  contains  more  than  a  small  amount  of  sodium,  is 
solid,  and  contains  one  or  more  compounds  of  the  two  elements. 
This  amalgam  is  often  used  instead  of  the  metal  sodium,  since 
the  dilution  and  combination  with  mercury  make  the  interactions 
of  the  metal  more  easily  controllable.  Sodium  is  used  in  the 
manufacture  of  sodium  peroxide  and  of  many  carbon  compounds 
which  are  used  as  drugs  and  dyes. 

Sodium  Chloride  NaCL  —  Common  salt  is  obtained  from  the 
salt  deposits  of  Stassfurt  and  Reichenhall  (near  Salzburg),  in 
Cheshire,  at  Syracuse  and  Warsaw  in  New  York,  at  Salina  in 
Kansas,  in  Utah,  California,  and  many  other  districts.  Natural 
brines  are  obtained  from  wells  in  various  parts  of  the  world.  Since 
the  salt  can  seldom  be  used  directly,  on  account  of  impurities  which 
it  contains,  it  is  purified  by  recrystallization  from  water.  Natural 
brines,  which  are  sometimes  dilute,  are  often  concentrated  by 
dripping  over  extensive  ricks  composed  of  twigs.  When  the  re- 
sulting brine  is  allowed  to  evaporate  slowly  by  the  help  of  the 
sun's  heat,  large  crystals,  sold  as  "  solar  salt,"  are  obtained.  By 
the  use  of  artificial  heat  and  stirring,  smaller  crystals  of  greater 
purity  can  be  secured.  In  northern  Russia,  the  brine  is  allowed  to 
freeze,  and  the  water  thus  removed  in  the  form  of  ice  (p.  134). 
Salt  intended  for  table  use  must  be  freed  from  the  traces  of  mag- 
nesium chloride  (q.v.)  present  in  the  original  brine  or  deposit,  for 
this  impurity  causes  it  to  absorb  moisture  more  vigorously  from 
the  air.  Addition  of  a  little  baking  soda  NaHCOa  remedies  the 
difficulty,  by  forming  the  insoluble  magnesium  carbonate.  The 
purest  salt  for  chemical  purposes  is  precipitated  from  a  saturated 
solution  of  salt  by  leading  into  it  hydrogen  chloride  gas.  Ex- 
planation of  this  effect  will  be  given  presently  (see  pp.  466-472). 

Common  salt  crystallizes  in  cubes,  the  faces  of  which  are  usually 
hollow.  The  crystals  decrepitate  (p.  449)  when  heated,  and  melt 
at  about  820°.  Common  salt  is  the  source  of  all  sodium  com- 
pounds, with  the  exception  of  the  nitrate.  From  it  come  also 


SODIUM  459 

most  of  the  chlorine  and  hydrogen  chloride  used  in  commerce. 
It  is  a  necessary  article  of  diet,  furnishing,  for  example,  the  hydro- 
chloric acid  in  the  gastric  juice  (p.  147). 

The  Hydroxide  and  Oxides.  —  Sodium  hydroxide  NaOH, 
called  also,  colloquially,  caustic  soda,  is  prepared  by  the  action  of 
slaked  lime  upon  sodium  carbonate,  but  mainly  by  the  electrolysis 
of  a  solution  of  sodium  chloride,  in  both  cases  precisely  as  is  potas- 
sium hydroxide  (p-.  446).  Sodium  hydroxide  is  a  highly  deliques- 
cent substance.  Its  general  chemical  properties  are  identical  with 
those  of  potassium  hydroxide.  It  is  used  in  the  manufacture  of 
soap,  in  the  preparation  of  paper  pulp,  and  in  many  other  chemical 
industries. 

Sodium  peroxide  Na202  is  made  by  heating  sodium  at  300-400° 
in  air  which  has  been  freed  from  carbon  dioxide.  The  sodium  is 
placed  on  trays  of  aluminium,  and  is  passed  into  the  furnace 
against  the  current  of  air.  In  this  way,  the  freshest  sodium  meets 
the  air  from  which  most  of  the  oxygen  has  been  removed,  and  the 
action  is  moderated.  Conversely,  the  almost  entirely  oxidized 
sodium  meets  the  freshest  air,  and  completion  of  the  oxidation  is 
thus  assured. 

This  oxide  is  the  sodium  salt  of  hydrogen  peroxide.  When 
thrown  into  water  it  decomposes  in  part,  in  consequence  of  the 
heat  developed,  giving  sodium  hydroxide  and  oxygen.  With  care- 
ful cooling,  however,  much  of  it  can  be  dissolved.  By  interaction 
with  acids  it  yields  hydrogen  peroxide  (p.  222).  Sodium  peroxide 
is  now  used  commercially  for  oxidizing  and  bleaching,  and,  in  the 
form  of  oxone  (p.  28),  as  a  source  of  oxygen. 
The  ordinary  sodium  oxide  Na2O  is  made 
in  the  same  way  as  is  potassium  oxide 
(p.  447). 

The  Nitrate  and  Nitrite.  —  The  occur- 
rence  and    purification   of   sodium   nitrate 

NaNO3  have  already  been  described  (p.  347).  Its  crystals  are 
of  rhombohedral  form  (Fig.  115).  This  salt  is  one  of  the  best  of 
fertilizers,  since  it  furnishes  to  plants  the  nitrogen  which  they 
require  in  a  very  easily  absorbed  form.  It  is  used  also  in  the 
manufacture  of  potassium  nitrate,  and  of  nitric  acid. 


460 


COLLEGE    CHEMISTRY 


Sodium  nitrite  NaN02  is  formed  by  heating  sodium  nitrate  with 
metallic  lead  and  recrystallizing  the  product  (p.  356). 

Manufacture  of  Sodium  Carbonate.  —  Natural  sodium  car- 
bonate is  found  in  Egypt  and  in  other  parts  of  the  world.  At 
Owen's  Lake,  California,  it  is  secured  by  solar  evaporation  of  the 
water.  The  sesquicarbonate  Na2C03,NaHC03,2H2O,  being  the 
least  soluble  of  the  carbonates  of  sodium,  is  the  one  deposited. 
Locally,  small  quantities  of  sodium  carbonate  are  still  made  by 
the  burning  of  sea-weed.  The  substance  is  manufactured  from 
sodium  chloride  in  two  ways,  namely  by  the  Le  Blanc  process  and 
by  the  Solvay  process.  In  1900,  however,  only  two  factories  used 
the  former  process. 

The  Le  Blanc  process  (1791)  involves  three  chemical  actions.  In 
the  first  place,  sodium  chloride  is  treated  with  an  equivalent 


FIG.  116. 

amount  of  sulphuric  acid  in  a  large  cast-iron  or  earthenware  pan. 
The  bisulphate  thus  produced  (cf.  p.  141),  together  with  the  un- 
changed sodium  chloride,  is  raked  out  on  to  the  hearth  of  a  rever- 
beratory*  furnace  (Fig.  116),  or  into  a  rotating,  inclined  iron 
cylinder,  and  heated  more  strongly  until  the  action  is  completed: 

NaCl  +  NaHS04  «=±  Na2S04  +  HC1 1 . 

*  So  called  because  the  heated  gases  from  the  fire  are  deflected  by  the  roof 
and  play  upon  the  materials  spread  on  the  bed  of  the  furnace. 


SODIUM  461 

The  product  of  this  treatment  is  called  salt  cake.  The  hydrogen 
chloride,  which  is  liberated  in  both  stages,  passes  through  towers 
containing  running  water  in  which  it  is  absorbed.  The  second  and 
third  actions  which  follow  are  conducted  in  one  operation.  They 
consist  in  the  reduction  of  the  sodium  sulphate  by  means  of 
powdered  coal  and  the  interaction  of  the  resulting  sulphide  of 
sodium  with  chalk  or  powdered  limestone,  leaving  finally  black  ash : 

Na2S04  +  2C  -»  Na2S  +  2CO2, 
Na2S  +  CaC03  ->  Na2CO3  +  CaS. 

Calcium  sulphide  is  not  very  soluble  in  water,  and  is  but  slowly 
hydrolyzed  by  it  (p.  273),  especially  when  calcium  hydroxide  is 
present.  The  sodium  carbonate  is  therefore  extracted  from  the 
black  ash  by  a  systematic  treatment  of  the  ash  with  water.  The 
ash  is  placed  in  a  series  of  vessels  at  different  levels,  and  a  stream 
of  water  (30-40°)  flows  from  one  vessel  to  another,  until,  when  it 
issues  from  the  last,  it  is  completely  saturated  with  sodium  car- 
bonate. When  the  material  in  the  first  of  the  vessels  has  been 
exhausted,  the  water  is  allowed  to  enter  the  second  vessel  directly, 
and  a  vessel  containing  fresh  black  ash  is  added  at  the  lower  end 
of  the  series.  In  this  way  the  most  nearly  exhausted  ash  comes  in 
contact  with  pure  water,  which  is  in  the  best  position  to  dissolve 
the  remaining  sodium  carbonate  rapidly,  while  the  fresh  black  ash 
encounters  a  solution  already  almost  at  the  point  of  saturation. 
The  commercial  survival  of  the  process  depends  upon  the  recovery 
of  the  sulphur  from  the  spent  black  ash,  and  of  the  hydrogen 
chloride. 

The  Solvay,  or  ammonia-soda  process  (1860),  has  now  displaced 
the  Le  Blanc  process.  It  differs  from  the  latter  by  involving 
almost  nothing  but  ionic  actions.  A  solution  of  salt,  containing 
ammonia  and  warmed  to  40°,  fills  a  tower  divided  by  a  number  of 
perforated  partitions.  Carbon  dioxide,  which  is  forced  in  below, 
makes  its  way  up  through  the  liquid.  The  ammonium  bicarbonate 
formed  by  its  action  undergoes  double  decomposition  with  the 
salt,  and  sodium  bicarbonate  which  is  precipitated  (sol'ty,  9.6  g. 
in  100  c.c.  Aq)  settles  upon  the  partitions: 

NaCl  +  NH4HC03  ±=>  NaHC03 1  +  NH4C1, 
or  HC03~  +  Na+  *±  NaHCO3 J . 


462  COLLEGE    CHEMISTRY 

The  solid  sodium  bicarbonate,  after  being  freed  from  the  liquid,  is 
heated  strongly  and  leaves  behind  sodium  carbonate: 

2NaHC03  «-  Na^COa  +  H20  |  +  CO2|  . 


The  carbon  dioxide  which  is  liberated  passes  through  the  operation 
once  more.  The  supply  of  carbon  dioxide  is  generated  in  lime- 
kilns of  special  form.  The  mother-liquor  from  the  sodium  bicar- 
bonate contains  ammonium  chloride.  This  is  decomposed  by 
heating  with  quicklime  from  the  kilns,  and  the  ammonia  which  is 
thus  obtained  is  available  for  the  treatment  of  another  batch. 

The  anhydrous  sodium  carbonate  (soda  ash  or  calcined  soda)  is 
recrystallized  from  water,  giving  the  decahydrate  Na2C03,10H2O, 
soda  crystals,  or  washing  soda.  The  bicarbonate  is  baking  soda. 

Properties  of  the  Carbonate  and  Bicarbonate.  —  The  com- 
mon form  of  sodium  carbonate  consists  of  large  monoclinic  crystals 
of  the  decahydrate  Na2C03,10H20.  This  substance  has  a  fairly 
high  aqueous  tension,  and  loses  nine  of  the  ten  molecules  of  water 
which  it  contains  when  it  is  exposed  in  an  open  vessel  (p.  96), 
leaving  the  monohydrate.  When  warmed  it  melts  at  35.2°,  giving 
a  solution  of  sodium  carbonate  in  water.  The  deposit  from  evap- 
oration, above  35.2°,  is  the  monohydrate  Na2C03,H20.  At  higher 
temperatures,  or  in  a  dry  atmosphere  (p.  96),  this  in  turn  can  be 
completely  dehydrated.  In  aqueous  solution,  sodium  carbonate 
is  hydrolyzed  (2.3  per  cent  in  0.1  N  solution  at  25°),  and  shows  a 
marked  alkaline  reaction  (p.  383)  .  The  compound  is  used  in  large 
amounts  for  the  manufacture  of  glass  and  soap,  and  in  the  soften- 
ing of  water,  and  is  applied  in  innumerable  ways  in  the  scientific 
industries  for  purposes  akin  to  cleansing. 

All  the  familiar  compounds  of  sodium,  excepting  sodium  nitrate 
and  the  peroxide,  are  made  by  the  treatment  of  sodium  carbonate 
or  sodium  hydroxide  with  acids. 

Sodium  bicarbonate  NaHC03  is  formed  in  the  Solvay  process 
(p.  461).  It  can  be  prepared  in  a  state  of  purity  by  passing  carbon 
dioxide  over  the  decahydrate  of  sodium  carbonate: 

Na2C03,10H20  +  C02  <=»  2NaHCO3  +  9H20. 

This  action  is  reversible  (cf.  p.  384),  and  sodium  bicarbonate  shows, 
even  in  the  cold,  an  appreciable  tension  of  carbon  dioxide.  The 


SODIUM  463 

aqueous  solution  of  the  pure  substance  is  neutral  to  phenolphthal- 
ei'n,  on  account  of  the  small  degree  of  ionization  of  the  ion  HCOs". 
Ordinarily,  however,  the  solution  is  alkaline,  on  account  of  the 
presence  of  the  carbonate,  which  is  hydrolyzed.  The  salt  is  used 
in  the  manufacture  of  baking  powder  and  in  medicine. 

Baking  Powders.  —  The  object  of  using  the  powder  is  to 
generate  carbon  dioxide  in  the  dough.  The  bubbles  are  retained 
because  of  the  presence  of  the  sticky  gluten,  a  protein  (p.  3). 
They  expand  when  the  dough  is  heated  in  baking,  and  give  to  the 
bread  its  porous  texture. 

Baking  soda,  alone,  will  give  off  carbon  dioxide,  but  the  sodium 
carbonate  which  it  leaves  behind  has  a  disagreeable  taste  and  acts 
upon  the  gluten  causing  a  yellow  color  and  unpleasant  smell.  It 
also  tends  to  neutralize  the  acid  in  the  gastric  juice  and  so  impedes 
digestion.  To  prevent  this  result,  sour  milk  (containing  lactic 
acid)  and  even  vinegar  are  added.  Usually,  however,  a  baking 
powder  containing  an  acid  substance  along  with  the  bicarbonate  is 
employed.  Potassium  bitartrate  (cream  of  tartar)  KHC4H406  (p. 
452)  is  most  commonly  employed,  although  alum  and  primary 
sodium  or  ammonium  orthophosphate  (p.  370)  are  also  used: 

NaHC03  -»  NaKC4H4O6  +  H2C03  -»  H20  +  C02. 


The  cream  of  tartar  has  the  advantages  that  it  is  somewhat  in- 
soluble and  does  not  act  noticeably  upon  the  soda  before  the 
mixing  of  the  dough  is  complete,  and  that  the  sodium-potassium 
tartrate  (Rochelle  Salts)  produced  is  not  harmful.  A  little  starch 
is  added  to  baking  powders  to  keep  the  particles  of  the  two  other 
ingredients  apart,  and  prevent  gradual  interaction  before  use. 

For  raising  bakers'  bread,  yeast  is  employed,  and  time  is  allowed 
for  the  propagation  of  the  yeast  and  its  action  upon  the  sugar  (p. 
406)  in  the  flour.  A  little  molasses  or  malt  extract  is  often  added, 
to  ensure  a  sufficient  supply  of  sugar. 

The  whites  of  eggs  cause  cake  to  rise,  largely  because  they  are 
whipped  before  use,  and  bubbles  of  air,  which  expand  when 
heated,  are  thus  introduced. 

Other  Salts  of  Sodium.  —  Anhydrous  sodium  sulphate  Na2S04 
(thenardite)  is  found  in  the  salt  layers.  The  same  salt  is  contained 


464  COLLEGE    CHEMISTRY 

in  mineral  waters,  such  as  those  of  Friedrichshall  and  Karlsbad. 
It  is  formed  in  connection  with  the  manufacture  of  nitric  acid  from 
sodium  nitrate.  It  is  used,  as  a  substitute  for  sodium  carbonate, 
in  making  inexpensive  glass. 

The  decahydrate  of  sodium  sulphate  Na2S04,10H2O  (Glauber's 
salt)  forms  large  monoclinic  crystals  which  give  up  all  their  water 
of  hydration  when  kept  in  an  open  vessel.  When  heated,  the 
crystals  melt  at  32.4°,  giving  the  sulphate  and  water.  For  the 
solubilities  of  the  hydrate  and  anhydrous  substance,  see  Fig.  59 
(p.  132). 

Sodium  thiosulphate  Na2S2O3,5H2O,  formerly  called  hyposulphite 
of  soda,  and  still  called  hypo  by  photographers,  is  made  by  boiling 
a  solution  of  sodium  sulphite  with  sulphur  (p.  290).  A  standard 
solution  (p.  257)  of  the  thiosulphate  is  used  in  determining 
quantities  of  free  iodine: 

2Na2S203  +  Is  ->  2NaI  + 


Colorless  sodium  tetrathionate  is  formed,  and  the  "end  point  " 
(consumption  of  all  the  iodine)  can  be  ascertained  by  the  starch 
test  (see  p.  480). 

When  heated,  dry  sodium  thiosulphate  first  loses  the  water  of 
hydration,  and  then  decomposes,  giving  sodium  sulphate,  which  is 
the  most  stable  oxygen-sulphur  compound  of  sodium  (cf.  p.  290) 
and  sodium  pentasulphide: 

4Na2S203  ->  3Na2S04  +  Na2S5. 

From  the  latter,  four  unit-weights  of  sulphur  can  be  driven  by 
stronger  heating.  Sodium  thiosulphate  is  used  for  fixing  negatives 
in  photography  (q.v.),  and  by  bleachers  as  an  antichlor. 

Sodium  hyposulphite  Na^SgC^  is  prepared  in  solution  by  the 
action  of  zinc  on  sodium  bisulphite  and  excess  of  sulphurous  acid: 

Zn  +  2NaHS03  +  H2S03  ->  Na2S204  +  ZnS03  +  2H20. 

The  solution  is  an  active  reducing  agent,  and  is  employed  largely 
by  dyers,  for  example  in  reducing  indigo  (insoluble)  to  indigo  white 
(soluble  in  an  alkaline  liquid),  in  preparing  the  vat  of  dye. 

Common  sodium  phosphate  is  a  dodecahydrate  of  the  secondary 
orthophosphate,  Na2HP04,12H2O.  It  is  made  by  neutralization 
of  phosphoric  acid  with  sodium  carbonate.  Its  properties  have 
already  been  discussed  (pp.  370-371). 


SODIUM  465 

Sodium  metaphosphate  NaPOs  is  formed  in  bead  tests  (p.  371). 

Sodium  tetraborate  Na2B4O7,10H20  (borax)  forms  large,  trans- 
parent prisms.  When  heated  it  loses  water,  and  leaves  the  easily 
fusible  anhydrous  salt  in  glassy  form.  Its  sources  have  already 
been  discussed  under  borates  (p.  432).  It  is  used  as  an  ingredient 
in  glazes  for  porcelain,  in  soldering,  for  bead  reactions  (p.  433)  and 
for  preserving  food. 

Sodium  disilicate  Na2Si205  (cf.  p.  428)  is  used  for  fireproofing 
wood  and  other  materials,  and  for  preserving  eggs.  Sand  which 
is  moistened  with  it  and  pressed  in  molds,  forms,  after  baking,  a 
serviceable  artificial  stone. 

For  sodium  cyanide,  see  p.  488. 

Properties  of  Sodium-ion  Na+:    Analytical  Reactions.  — 

Sodium-ion  is  a  colorless  ionic  material  which  unites  with  all 
negative  ions.  Practically  all  the  salts  so  formed  are  soluble  in 
water.  The  only  ones  which  can  be  precipitated  are  sodium  fluo- 
silicate  Na2SiF6,  made  by  the  addition  of  hydrofluosilicic  acid  to  a 
strong  solution  of  a  sodium  salt,  and  sodium-hydrogen  pyroanti- 
moniate  Na2H2Sb207,  made  by  similar  addition  of  the  corresponding 
potassium  salt.  All  compounds  of  sodium  confer  a  yellow  color 
on  the  Bunsen  flame,  but  this  test  is  so  delicate  that  it  is  shown  by 
the  traces  of  sodium  contained  in  almost  all  substances. 

Lithium.  —  Lithium  occurs  in  lepidolite  (a  lithia  mica),  in 
amblygonite,  and  in  other  rare  minerals.  Traces  of  compounds  of 
the  element  are  found  widely  diffused  in  the  soil,  and  are  taken  up 
by  plants^  particularly  tobacco  and  beets,  in  the  ashes  of  which  the 
element  may  be  detected  spectroscopically. 

The  metal  is  liberated  by  electrolysis  of  the  fused  chloride.  The 
specific  gravity  of  the  free  element  (0.53)  is  lower  than  that  of  any 
other  metal.  Lithium  not  only  floats  upon  water,  but  also  in  the 
petroleum  in  which  it  is  preserved. 

The  metal  behaves  towards  water  and  oxygen  like  sodium  (p.  50). 
It  unites  directly  and  vigorously  with  hydrogen  (LiH),  nitrogen 
(Li3N),  and  oxygen  (Li2O),  forming  stable  compounds.  The  rel- 
ative insolubility  (see  Table)  of  the  hydroxide  LiOH,  the  car- 
bonate LioCOa,  and  the  phosphate  Li3P04,2H20  is  in  sharp 
contrast  to  the  easy  solubility  of  the  corresponding  compounds  of 


466  COLLEGE    CHEMISTRY 

the  other  alkali-metals,  and  links  lithium  with  magnesium.  The 
compounds  of  lithium  give  a  bright-red  color  to  the  Bunsen  flame. 
A  bright-red  and  a  somewhat  less  bright  orange  line  are  seen  in  the 
spectrum.  The  carbonate  is  used  in  medicine. 

IONIC  EQUILIBRIUM,  CONSIDERED  QUANTITATIVELY 

In  view  of  the  predominance  of  ionic  actions  in  the  chemistry  of 
the  metals,  and  of  the  determinative  effect  of  ionic  equilibria  on 
many  actions,  it  is  essential  that  we  should  be  prepared  in  future 
for  a  more  exact  consideration  of  these  phenomena  than  we  have 
hitherto  attempted.  The  whole  basis  for  this  exact  consideration 
has  already  been  supplied,  and  only  more  specific  application  of  the 
principles  is  demanded.  The  basis  referred  to,  which  should  now 
be  re-read  as  a  preliminary  to  what  follows,  is  contained  in,  (1)  the 
discussion  of  chemical  equilibrium  in  general  (pp.  177-190),  (2)  the 
application  of  the  same  principles  to  ionic  equilibrium  (p.  238), 
and  (3)  the  illustration  of  this  application  in  the  case  of  cupric 
bromide  (pp.  246-251). 

Excess  of  One  Ion.  —  In  the  case  of  cupric  bromide,  we  showed 
that  increasing  the  concentration  of  the  bromide  ions  displaced  the 
equilibrium  by  favoring  the  union  of  the  ions  to  form  molecular 
cupric  bromide:  2Br~  +  Cu++  — >  CuBr2.  This  we  speak  of  as  a 
repression  of  the  ionization  of  the  cupric  bromide.  Now,  if  the  sub- 
stance is  a  slightly  ionized  one,  like  a  weak  acid  or  a  weak  base,  the 
repression  of  the  ionization  through  the  formation  of  molecules  in 
this  way  may  remove  so  many  of  that  one  of  the  ions  which  is  not 
present  in  excess  (corresponding  to  the  Cu++  in  the  foregping  illus- 
tration), that  the  mixture  will  no  longer  respond  to  tests  for  the 
ion  so  removed.  This  is  an  interesting  and  very  common  case. 
The  behavior  of  acetic  acid,  a  weak,  slightly  ionized  acid,  will  serve 
as  an  illustration. 

In  normal  solution  (60  g.  in  1  1.)  acetic  acid  is  only  0.004  ionized 
(p.  241),  so  that,  in  the  equation  for  the  equilibrium, 

(0.996)  HC2H302  ±5  H+  (0.004)  +  C2H3O2-  (0.004), 

the  relative  proportions  are  as  shown  by  the  numbers  in  parenthe- 
sis. If  the  whole  of  the  acid  (60  g.)  were  ionized,  there  would  be 
1  g.  of  hydrogen-ion  per  liter.  Yet,  even  in  the  much  smaller 


IONIC   EQUILIBRIUM,    CONSIDERED    QUANTITATIVELY      467 

concentration  actually  present  (0.004  g.  per  liter),  the  acid  taste  of 
the  H+  and  its  effect  upon  indicators  can  be  distinctly  recognized. 
If,  now,  solid  sodium  acetate  is  dissolved  in  the  solution,  the  liquid  no 
longer  gives  an  add  reaction  with  one  of  the  less  delicate  indicators, 
like  methyl  orange  (p.  258).  The  explanation  is  simple.  Sodium 
acetate  is  highly  ionized.  It  gives,  therefore,  a  large  concentra- 
tion of  acetate-ion  to  a  liquid  formerly  containing  very  little. 
This  causes  a  greatly  increased  union  of  the  H+  ions  and  C2H302~ 
ions  to  occur,  and  the  former,  being  already  very  few  in  number, 
disappear  almost  entirely.  Hence  the  solution  becomes,  to  all 
intents  and  purposes,  neutral.  There  is  no  less  acetic  acid  present 
than  before,  but  the  concentration  of  hydrogen-ion  is  very  much 
smaller. 

Formulation  and  Quantitative  Treatment  of  the  Case  of 

Excess  of  One  Ion.  —  If  the  semi-mathematical  mode  of  formu- 
lating an  equilibrium  (p.  184),  as  applied  to  the  case  of  an  ionogen 
(p.  238),  be  employed  here,  the  foregoing  general  statements  may 
be  made  more  precise  and  the  conclusions  clearer.  If  [H+]  and 
[C2H302~]  represent  the  molecular  concentrations  of  hydrogen-ion 
and  acetate-ion,  respectively,  and  [HC2H302]  that  of  the  acetic 
molecules  at  equilibrium,  then: 

[H+i  x  [cjfron  _  „ 

[HC»B*)J 

The  value  of  K  is  constant,  whether  the  strength  of  the  solution  of 
acetic  acid  is  great  or  small,  and  even  when  another  substance  with 
a  common  ion  is  present.  In  the  latter  case,  [C2H302~]  and  [H+] 
stand  for  the  whole  concentrations  of  each  of  these  ionic  substances 
from  both  sources. 

Now,  in  normal  acetic  acid  [H+]  =  0.004,  [C2H302~]  =  0.004  (for 
the  number  of  each  kind  of  ions  is  th^same),  and  [HC2H302]  = 
0.996,  practically  1.     Substituting  in  the  formula: 
0004X0004 


When,  however,  sodium  acetate  is  dissolved  in  the  liquid  until  the 
solution  is  normal  in  respect  to  this  substance  also,  the  following 
additional  equilibrium  has  to  be  considered: 

(0.47)  NaC2H302  <=±  Na+  (0.53)  +  C2H302-  (0.53). 


468  COLLEGE    CHEMISTRY 

The  concentration  of  acetate-ion  from  this  source  is  0.53,  so  that,  in 
the  mixture  of  acid  and  salt,  the  concentration  of  acetate-ion 
[C2H302~]  will  be  0.53  +  0.004  =  0.534,  or  nearly  134  times  larger 
than  in  the  acid  alone.  Hence,  in  order  that  the  product  [H+]  X 
[C2H3O2~]  may  recover,  as  it  must,  a  value  much  nearer  to  the  old 
one,  [H+]  must  be  diminished  to  something  like  T^¥  of  its  former 
magnitude.  That  is,  [H+]  will  become  equal  to  about  0,00003, 

0.00003  X  0.534 


the  rest  of  the  hydrogen-ion  uniting  with  a  corresponding  amount 
of  the  acetate-ion  to  form  molecular  acetic  acid.  The  effect  of 
adding  this  amount  of  sodium  acetate  therefore  is,  as  we  have 
seen,  to  reduce  the  concentration  of  the  hydrogen-ion  below  the 
amount  which  can  be  detected  by  use  of  an  indicator  like  methyl 
orange. 

This  effect  is  of  course  reciprocal,  and  the  ionization  of  the 
sodium  acetate  will  be  reduced  also.  But  the  acetate-ion  furnished 
by  the  acetic  acid  is  relatively  so  small  in  amount  (0.00003  against 
0.53)  that  the  effect  it  produces  on  the  ionization  of  the  salt  is 
imperceptible. 

It  will  be  noted  that  the  acetate-ion  and  hydrogen-ion  disappear 
in  equivalent  quantities,  for  they  unite.  There  is,  however,  so 
much  of  the  former  that  the  loss  it  sustains  goes  unremarked,  while 
there  is  so  little  of  the  latter  that  almost  none  of  it  remains.  When 
substances  of  more  nearly  equal  degrees  of  ionization  are  used, 
both  effects  are  equally  inconspicuous.  Thus,  sodium  chloride  and 
hydrogen  chloride  in  normal  solutions  yield  approximately  equal 
concentrations  of  chloride-ion  (0.784  and  0.66).  Hence,  if  one 
mole  of  sodium  chloride  were  to  be  dissolved  in  the  portion  of  water 
already  containing  one  mole  of  hydrogen  chloride,  the  concentra- 
tion of  the  chloride-ion,  at  a  very  rough  estimate,  would  be  nearly 
doubled.  If  this  doubling  of  the  concentration  of  chloride-ion 
almost  halved  that  of  the  hydrogen-ion  (0.784),  in  order  that  the 
expression  [Cl~]  X  [H+]  -5-  [HC1]  might  remain  constant,  the 
concentration  of  the  hydrogen-ion  would  still  be  about  0.400  and 
therefore  100  times  as  great  as  in  molar  acetic  acid.  It  is  thus 
altogether  impossible  to  reduce  the  concentration  of  the  hydrogen- 
ion  given  by  an  active  acid  like  hydrochloric  acid  below  the  limit 


IONIC   EQUILIBRIUM,    CONSIDERED   QUANTITATIVELY      469 

at  which  indicators  are  affected,  for  there  is  no  way  of  introducing 
the  enormous  concentration  of  the  other  ion  which  the  theory 
demands. 

With  more  crude  means  of  observation  than  indicators  afford, 
effects  like  this  last  may  sometimes  be  rendered  visible.  This  was 
the  case  with  cupric  bromide  solution,  to  which  potassium  bromide 
was  added  (p.  250).  The  blue  of  the  cupric-ion  disappeared  from 
view,  while  much  cupric-ion  was  still  present,  because  the  brown 
color  of  the  molecular  cupric  bromide  covered  it  up  completely. 

Special  Case  of  Saturated  Solutions.  —  The  commonest  as 
well  as  the  most  interesting  application  of  the  conceptions  de- 
veloped above  is  met  with  in  connection  with  saturated  solutions, 
especially  those  of  relatively  insoluble  substances. 

The  situation  in  a  system  consisting  of  the  saturated  solution 
and  excess  of  the  solute  has  been  discussed  already  (read  p.  127). 
In  the  case  of  potassium  chlorate,  for  example,  we  have  the1  follow- 
ing scheme  of  equilibria: 

KC103  (solid)  <=±  KC103  (dslvd)  <=>  K+  +  C103~ 

Solution  of  the  solid  is  promoted  by  the  solution  pressure  of  the 
molecules,  while  it  is  opposed  by  the  osmotic  pressure  of  the  dis- 
solved substance,  and  the  solution  is  saturated  when  these  tenden- 
cies produce  equal  effects  (p.  128).  Now  it  must  be  noted  that  the. 
tendency  directly  opposed  to  the  solution  pressure  is  the  partial 
osmotic  pressure  of  the  dissolved  molecules  alone.  The  chief  con- 
tents of  the  solution,  the  molecules  and  two  kinds  of  ions  of  the  salt, 
and  any  foreign  material  that  may  be  present,  are  like  a  mixture  of 
gases,  and  the  principle  of  partial  pressure  (p.  72)  is  to  be  applied. 
The  ions  and  the  foreign  material  do  not  deposit  themselves  upon 
the  solid,  and  take,  therefore,  no  part  directly  in  the  equilibrium 
which  controls  solubility.  In  respect  to  this  the  ions  are  them- 
selves foreign  substances.  Hence  the  conclusion  may  be  stated 
that,  in  solutions  saturated  at  a  given  temperature  by  a  given  solute, 
the  concentration  of  the  dissolved  molecules  of  the  solute  consid- 
ered by  themselves  will  be  constant  whatever  other  substances  may  be 
present. 

The  total  " solubility"  of  a  substance,  as  we  have  used  the  term 
hitherto,  is  made  up  of  a  molecular  and  an  ionic  part.  The  latter, 


470  COLLEGE    CHEMISTRY 

as  we  shall  presently  see,  is  not  constant  when  a  foreign  substance 
containing  a  common  ion  is  already  in  the  liquid.  Since  the  treat- 
ment of  the  subject  requires  us  now  to  distinguish  between  the  two 
portions  of  the  solute,  a  diagram  (Fig.  117)  will 
assist  in  emphasizing  the  distinction.  The  ma- 
terial at  the  bottom  is  the  salt.  The  mole- 
cules and  ions  are  to  be  thought  of  as  being 
mixed  and  as  being  present  in  numbers  repre- 


»K++mCIO, 


]  | 


KCIO. 


sented  by  the  factors  n  and  m.  Since  no  foreign 
body  is  present,  the  two  ions  in  this  case  are 
equal  in  number. 

When  we  now  apply  these  ideas  to  the  mathe- 
matical expression  of  the  relation: 

[K+]  X  [ClOT] 

^-  [KC103] 

we  perceive  that,  in  a  saturated  solution,  [KC1O3],  the  concentra- 
tion of  the  molecules,  is  constant.  Transposing,  we  have 

[K+]  X  [C10S-]  =  #[KCIO3]  =  Kf. 

Hence  the  relation  leads  to  the  important  conclusion  that,  in  a 
saturated  solution,  the  product  of  the  molar  concentrations  of  the 
ions  is  constant.*  This  product  is  called  the  ion-product  constant 
for  the  substance.  The  law  of  the  constancy  of  the  ion-product 
in  a  saturated  solution  is  one  of  the  most  useful  of  the  principles  of 
chemistry.  It  enables  us  to  explain  all  the  varied  phenomena  of 
precipitation  and  of  the  solution  of  precipitates  in  a  consistent 
manner.  These  applications  of  the  principle  will  be  explained  in 
the  next  chapter.  One  curious  kind  of  precipitation  will  be  de- 
scribed here,  however,  as  an  illustration  of  the  use  of  the  principle. 

Illustration  of  the  Principle  of  Ion-Product  Constancy.  — 

When,  to  a  saturated  solution  of  one  of  the  less  soluble  salts,  a 

*  The  principle  of  constant  concentration  of  dissolved  molecules,  stated 
above,  has  been  shown  to  express  the  facts  very  inaccurately.  Now  the 
principle  of  the  constancy  of  the  ratio  of  the  ion-product  to  the  concentration 
of  the  molecules  is  also  inaccurate  in  the  case  of  highly  ionized  substances, 
yet  in  such  a  way  that  the  two  errors  neutralize  one  another.  'Thus,  the 
principle  of  ion-product  constancy  here  given  is  in  itself  fairly  exact. 


IONIC   EQUILIBRIUM,    CONSIDERED   QUANTITATIVELY      471 

strong  solution  of  a  very  soluble  salt  having  one  ion  in  common 
with  the  first  salt  is  added,  precipitation  of  the  first  salt  frequently 
takes  place.  This  happens,  for  example,  with  a  saturated  solution 
of  potassium  chlorate,  which  is  not  very  soluble  (molar  solubility 
0.52,  see  Table) .  The  concentrations  [K+]  and  [C103~]  being  small, 
one  may  easily  increase  the  value  for  one  of  the  ions,  say  [C103~], 
fivefold,  by  adding  a  chlorate  which  is  sufficiently  soluble.  To 
preserve  the  value  of  the  product  [K+]  X  [C103~],  the  value  of 
[K+]  will  then  have  to  be  diminished  at  once  to  one-fifth  of  its 
former  value.  This  can  occur  only  by  union  of  the  ionic  material 
it  represents  with  an  equivalent  amount  of  that  for  which  [C103~] 
stands.  The  molecular  material  so  produced  will  thus  tend  at  first 
to  swell  the  value  of  [KC103].  But  the  value  of  [KC1O3]  cannot  be 
increased,  for  the  solution  is  already  saturated  with  molecules,  so 
that  the  new  supply  of  molecules,  or  others  in  equal  numbers,  will 
be  precipitated.  Hence  the  ionic  part  of  the  dissolved  substance 
may  be  diminished,  the  equilibria  (p.  469)  may  be  partially  re- 
versed, and  we  may  actually  precipitate  a  part  of  the  dissolved 
material  without  introducing  any  substance,  which,  in  the  ordinary 
sense,  can  interact  with  it. 

In  point  of  fact,  when,  to  a  saturated  solution  of  potassium 
chlorate  there  is  added  a  saturated  solution  of  potassium  chloride 
KC1  (molar  solubility,  3.9)  or  of  sodium  chlorate  NaC103  (molar 
solubility,  6.4),  a  precipitate  of  potassium  chlorate  is  thrown  down. 
These  two  salts,  each  containing  one  of  the  ions  of  KC1O3,  and  being 
much  more  soluble  than  the  latter  (see  Table),  increase  the  con- 
centration of  one  ion  and  cause  the  precipitation  in  the  fashion 
just  explained. 

The  product  of  the  concentrations  of  the  ions,  for  example  [K+] 
X  [C103~],  is  called  also  the  solubility  product,  because  these  two 
values  jointly  determine  the  magnitude  of  the  solubility  of  the 
substance.  The  solubility  of  the  molecules  is  irreducible,  but  the 
ionic  part  of  the  dissolved  material  may  become  vanishingly  small 
if  the  value  of  either  [X+]  or  [Y~]  is  very  minute.  The  ionic  part 
of  any  particular  substance  is  made  up  of  the  smaller  of  the  two 
concentrations  of  the  ionic  substances  which  it  yields,  plus  an 
equivalent  amount,  and  no  more,  of  the  concentration  of  the  other 
ion.  The  rest  of  the  other  ionic  substance  is  part  of  the  solubility 
of  some  other  component. 


472  COLLEGE    CHEMISTRY 

Other  Illustrations.  —  The  precipitation  of  sodium  chloride 
from  a  saturated  solution,  by  the  introduction  of  gaseous  hydrogen 
chloride  (p.  458),  is  to  be  explained  in  the  same  manner.  The 
equilibria  : 

NaCl  (solid)  <=»  NaCl  (dslvd)  ^  Na+  +  Cl- 
are reversed  by  the  introduction  of  additional  Cl~  from  the  very 
soluble,  and  highly  ionized  HC1. 

A  steady  stream  of  hydrogen  chloride  is  often  obtained  by  drop- 
ping concentrated  sulphuric  acid  into  saturated  hydrochloric  acid: 

H+  +  Cr  <=±  HC1  (dslvd)  ^  HC1  (gas). 

The  effect  is  due  in  part  to  repression  of  the  ionization  of  the  hydro- 
gen-chloride and  elimination  of  molecules  of  the  gas  from  the  water 
which  is  already  saturated  with  molecules  of  the  same  kind. 

The  formation  of  potassium  hydroxide  (p.  446)  ceases  when  a 
certain  concentration  has  been  reached.  This  occurs  because  the 
concentration  of  OH~,  which  rapidly  increases,  is  a  factor  in  the 
solubility  product  of  calcium  hydroxide,  [Ca++]  X  [OH~]2.  With 
much  OH~,  little  Ca"1"1"  is  required  to  give  the  constant,  numerical 
value  of  the  product.  When  the  concentration  [Ca++]  from  the 
hydroxide  has  become  about  as  small  as  that  from  the  carbonate, 
the  motive  for  the  interaction  has  been  removed.  This  principle 
is  thus  as  important  in  industrial  operations  as  it  is  in  analytical 
and  other  laboratory  experimentation. 

Exercises.  —  1.  The  vapor  density  of  sodium  peroxide  has  not 
been  determined.  Why  is  the  formula  Na2O2  assigned  to  it? 

2.  Construct  a  scheme  of  equilibria  (p.  271)  showing  the  hy- 
drolysis of  calcium  sulphide.     Why  does  the  presence  of  calcium 
hydroxide  diminish  the  tendency  to  hydrolysis  (p.  461)? 

3.  What  will  be  the  effect  of  adding  a  concentrated  solution  of 
silver  nitrate  to  a  saturated  solution  of  silver  sulphate  (see  Table 
of  solubilities)? 

4.  Although  a  20  per  cent  solution  of  soap  can  easily  be  made, 
a  0.5  per  cent  solution  can  be  salted  out  (p.  417).     How  does  this 
fact  show  that  salting  out  is  not  an  operation  like  the  precipitations 
just  discussed? 


CHAPTER  XXXVI 
THE  METALLIC  ELEMENTS  OF  THE  ALKALINE  EARTHS 

The  Chemical  Relations  of  the  Elements.  —  The  familiar 
metals  of  this  group,  calcium  (Ca,  at.  wt.  40.1),  strontium  (Sr,  at. 
wt.  87.6),  and  barium  (Ba,  at.  wt.  137.4),  constitute  a  typical 
chemical  family,  both  in  the  qualitative  resemblance  to  one  an- 
other of  the  elements  and  of  the  corresponding  compounds,  and 
in  the  quantitative  variation  in  the  properties  with  increasing 
atomic  weight.  The  metals  themselves  displace  hydrogen  vigor- 
ously from  cold  water,  giving  hydroxides.  The  solutions  of  these 
hydroxides,  although  dilute,  on  account  of  a  rather  small  solu- 
bility, are  strongly  alkaline  in  reaction.  The  high  degree  of  ion- 
ization  of  the  hydroxides  recalls  the  hydroxides  of  the  metals  of 
the  alkalies,  and  their  relative  insolubility  the  hydroxides  of  the 
" earths"  (q.v.). 

In  all  their  compounds,  calcium,  strontium,  and  barium  are 
bivalent.  The  hydroxides  are  formed  by  union  of  the  oxides  with 
water,  and  are  progressively  less  easy  to  decompose  by  heating, 
barium  hydroxide  being  the  hardest.  The  carbonates,  when 
heated,  yield  the  oxide  of  the  metal  and  carbon  dioxide,  barium 
carbonate  being  the  most  difficult  to  decompose.  The  nitrates, 
when  heated  moderately,  give  the  nitrites,  but  the  latter  are 
broken  up  by  further  heating  and  yield  the  oxide  of  the  metal, 
and  nitrogen  tetroxide.  In .  these  and  other  respects  the  com- 
pounds of  the  metals  of  the  alkaline  earths  resemble  those  of  the 
heavy  metals  and  differ  from  those  of  the  metals  of  the  alkalies. 
Barium  approaches  the  latter  most  nearly. 

The  table  of  solubilities  (q.v.)  shows  that  the  chlorides  and 
nitrates  of  calcium,  strontium,  and  barium  are  all  soluble  in 
water,  the  solubility  diminishing  in  the  order  given.  The  sulphates 
and  hydroxides  cover  a  wide  range  from  slight  solubility  to  ex- 
treme insolubility.  Of  the  sulphates,  2100,  110,  and  2.3  parts, 
respectively,  dissolve  in  one  million  parts  of  water.  In  the  case 

473 


474  COLLEGE    CHEMISTRY 

of  the  hydroxides  the  order  of  magnitude  is  reversed,  and  the  cor- 
responding numbers  are  200,  630,  and  2200.  The  carbonates  are 
almost  as  insoluble  as  is  barium  sulphate.  Radium  (Ra,  at.  wt. 
226)  belongs  to  this  family  (see  under  Uranium). 

CALCIUM  Ca 

Occurrence.  —  The  fluoride,  and  the  various  forms  of  the  car- 
bonate, sulphate,  and  phosphate,  which  are  found  In  nature,  are 
described  below.  As  silicate,  calcium  occurs,  along  with  other 
metals,  in  many  minerals  and  rocks.  Compounds  of  the  element 
are  found  also  in  plants,  and  in  the  bones  and  shells  of  animals. 

The  Metal.  —  Calcium  is  made  by  electrolysis  of  the  molten 
chloride.  A  hollow  cylinder  made  of  blocks  of  carbon  bolted 
together  and  open  above,  forms  the  anode.  A  rod  of  copper 
hanging  so  that  its  end  dips  into  the  melt  forms  the  cathode. 
The  melting  of  the  anhydrous  calcium  chloride  with  which  the 
cylinder  is  filled  is  started  by  means  of  a  thin  rod  of  carbon  laid 
across  from  the  anode  to  the  cathode.  When  the  heat  generated 
by  the  passage  of  the  current  through  this  highly  resisting  medium 
has  melted  a  sufficient  amount  of  the  salt,  the  rod  is  removed,  and 
the  resistance  of  the  fused  material  suffices  to  maintain  the  tem- 
perature. The  calcium  rises  round  the  cathode  and  collects  on 
the  surface  of  the  bath.  By  slowly  elevating  the  copper  cathode, 
the  calcium,  which  adheres  to  it,  may  be  drawn  out  of  the  fused 
mass  in  the  form  of  a  gradually  lengthening,  irregular  rod.  The 
rod  of  calcium  is  kept  constantly  in  contact  with  the  metal 
which  accumulates  on  the  surface,  and  thus  forms  one  of  the 
electrodes. 

Calcium  is  a  silver-white,  crystalline  metal  (m.-p.  800°,  sp.  gr. 
1.55)  which  is  a  little  harder  than  lead,  and  can  be  cut,  drawn, 
and  rolled.  It  interacts  rapidly  with  water.  When  heated  it 
unites  vigorously  with  hydrogen,  oxygen,  the  halogens,  and 
nitrogen.  On  this  account  it  is  used  in  producing  a  high  degree 
of  evacuation.  It  burns  in  the  air,  giving  a  mixture  of  the  oxide 
and  nitride  Ca3N2.  The  presence  of  the  latter  may  be  shown  by 
the  liberation  of  ammonia  when  water  is  added  to  the  residue: 

Ca3N2  +  6H20  -» 3Ca(OH)2  +  2NH3. 


CALCIUM  475 

A  white  crystalline  hydride  CaH2  is  formed  by  direct  union  of 
the  constituents.  It  is  known  in  commerce  as  hydrolyte.  It  is 
an  expensive,  but  portable  source  of  hydrogen  for  filling  war 
balloons : 

CaH2  +  2H20  -H>  Ca(OH)2  +  2H2. 

Calcium  Chloride  CaCl2*  —  This  salt,  for  which  there  is  no 
extensive  commercial  application,  is  formed  as  a  by-product  in 
many  industrial  operations.  Thus,  it  is  a  by-product  of  the 
Solvay  soda  process  (p.  461).  By  evaporation  of  any  solution, 
the  hexahydrate  CaCl2,6H20  is  obtained  in  large,  deliquescent, 
six-sided  prisms.  On  account  of  the  great  concentration  of  a 
saturated  solution  of  this  compound,  the  solid  and  solution  do 
not  reach  a  condition  of  equilibrium  with  ice  (cf.  p.  134)  until  the 
temperature  has  fallen  below  —50°.  The  Solvay  process  brine 
(p.  462)  when  mixed  with  ice,  gives,  therefore,  a  very  efficient 
freezing  mixture.  On  account  of  its  deliquescent  character,  the 
solid  salt  is  sprinkled  on  roads  to  lay  the  dust. 

Calcium  chloride,  partly  dehydrated  by  heating,  CaCl2,2H2O, 
forms  a  porous  mass  which  is  used  in  chemical  laboratories  for 
drying  gases  and  liquids.  When  complete  dehydration  is  at- 
tempted, the  salt  interacts  with  the  water,  giving  some  calcium 
oxide. 

Calcium  chloride  forms  compounds,  not  only  with  water,  but 
also  with  ammonia  (CaCl2,8NH3)  and  with  alcohol.  For  drying 
these  substances,  therefore,  quicklime  is  employed. 

Calcium  Fluoride  CaFQ.  —  This  compound  occurs  in  nature  as 
fluorite  or  fluor-spar  CaF2.  It  crystallizes  in  cubes,  is  insoluble  in 
water,  and  when  pure  is  colorless.  Natural  specimens  often 
possess  a  green  tint  or  show  a  violet  fluorescence.  It  is  formed  as  a 
precipitate  when  a  soluble  fluoride  is  added  to  a  solution  of  a  salt 
of  calcium. 

Fluorite  is  used  in  the  etching  of  glass,  as  the  source  of  the 
hydrogen  fluoride  (p.  205).  It  is  easily  fusible,  as  its  name  indi- 
cates (Lat.  fluere,  to  flow),  and  is  employed  in  metallurgical 
operations  as  a  flux  (p.  438),  for  lowering  the  melting-point  (or 
freezing-point,  which  is  the  same  thing,  cf.  p.  134)  of  the  slag  (p. 
438),  and  so  facilitating  the  separation  of  the  latter  from  the  metal. 


476  COLLEGE    CHEMISTRY 

Calcium  Carbonate  CaCO3-  —  This  compound  is  found  very 
plentifully  in  nature.  Limestone  is  a  compact,  indistinctly 
crystalline  variety,  while  marble  is  a  distinctly  crystalline  form. 
Chalk*  is  a  deposit  consisting  of  the  calcareous  parts  of  minute 
organisms.  Egg-shells,  oyster-shells,  coral,  and  pearls  are  other 
varieties  of  organic  origin.f  Calcite  and  Iceland  spar  (Ger. 
spalten,  to  split)  are  pure  crystallized  calcium  carbonate.  The 
former  occurs  in  flat  rhombohedrons,  or  in  pointed,  six-sided 
crystals  (Fig.  43,  p.  83)  (scalenohedrons)  of  "  dog-tooth"  spar,  be- 
longing to  the  same  system. 

When  heated,  calcium  carbonate  dissociates,  giving  carbon  dioxide 
and  quicklime: 


At  ordinary  temperatures  the  decomposition  is  imperceptible. 
On  the  contrary,  atmospheric  carbon  dioxide,  in  spite  of  its  very 
low  partial  pressure,  combines  with  quicklime,  giving  "air- 
slaked"  lime.  As  the  temperature  rises,  however,  the  tension  of 
carbon  dioxide  coming  from  the  carbonate  increases,  and  has  a 
fixed  value  for  each  temperature.  If  it  is  continuously  allowed  to 
escape,  so  that  the  maximum  pressure  is  not  reached,  the  whole 
of  the  salt  eventually  decomposes.  At  700°  the  pressure  is  only 
25  mm.,  at  900°  it  reaches  an  atmosphere,  and  at  950°  two  atmos- 
pheres. The  phenomenon  is  precisely  similar  to  the  dissociation 
of  a  hydrate  (p.  96)  and  to  the  evaporation  of  a  liquid  (p.  88). 

Limestone  is  soluble  in  water  containing  carbonic  acid,  giving 
calcium  bicarbonate  (p.  384,  also  see  p.  489).  By  solution  of 
limestone,  caves  are  often  formed.  Conversely,  subterranean 
water  containing  the  bicarbonate,  when  it  reaches  such  a  cavern, 
loses  carbon  dioxide  and  deposits  calcium  carbonate  as  stalactites 
or  columns  hanging  from  the  ceiling.  The  drippings  form  stalag- 
mites on  the  floors. 

Limestone  is  used  in  the  manufacture  of  quicklime  (q.v.)  and  of 
glass.  It  is  employed  largely  as  a  flux  in  metallurgy,  when  min- 
erals rich  in  silica  are  brought  into  fusible  form  by  the  production 
of  calcium  silicate  CaSiOa.  Large  amounts  also  find  application 
as  building-stone. 

*  Blackboard  "  crayon"  is  usually  made  of  gypsum  and  not  of  chalk. 
t  The  hard  coverings  of  Crustacea  and  insects  are  not  made  of  this  sub- 
stance, but  of  an  organic  material  called  chitin. 


CALCIUM 


477 


CAP  BON 
~~  oioxtoe 


Calcium  Oxide  and  Hydroxide.  —  Pure  oxide  of  calcium  CaO 
(quicklime)  may  be  made  by  ignition  of  pure  marble  or  calcite. 
For  commercial  purposes  limestone  is  converted  into  quicklime 
in  kilns  (Fig.  118).  The  flames  and  heated  gases  from  the  fire 
pass  between  the  pieces  of  limestone,  and  the  carbon  dioxide 
liberated  is  carried  off  by  the  draft. 
When  the  gas  is  to  be  used  in  the 
Solvay  process  or  in  the  refining  of 
sugar,  coke  (smokeless),  instead  of 
coal,  is  employed  as  the  fuel.  As 
low  a  temperature  as  possible  is 
used.  A  high  temperature  causes 
impurities  in  the  limestone  (e.g., 
clay)  to  interact  with  the  quick- 
lime, giving  fusible  silicates,  which 
fill  the  pores  and  interfere  with 
the  subsequent  slaking  with  water. 
Since  the  change  is  reversible,  if  the 
gas  lingers  in  the  kiln  (at  760  mm. 
pressure),  a  temperature  over  900° 
is  required  to  drive  the  action  forward  (p.  476).  Hence,  a  good 
draft,  which  removes  the  gas  as  fast  as  it  is  formed,  permits  the 
use  of  a  lower  temperature. 

Pure  calcium  oxide  is  a  white,  porous  solid.  It  is  barely  fusible 
in  the  oxyhydrogen  flame,  but  may  be  melted  and  boiled  in  the 
electric  arc.  It  is  not  reducible  by  sodium,  or  by  carbon  excepting 
at  the  temperature  of  the  electric  furnace. 

When  water  is  poured  upon  quicklime,  it  is  first  absorbed  into 
the  pores  mechanically,  and  then  unites  chemically  to  form  calcium 
hydroxide  Ca(OH)2: 

H20±=FCa(OH)2. 


Fia.  118. 


The  product  is  a  bulky  powder.  Much  heat  is  evolved,  and  part 
of  the  water  is  turned  into  steam.  The  change  is  reversible,  and 
at  a  high  temperature  the  hydroxide  can  be  dehydrated. 

Calcium  hydroxide  is  slightly  soluble  in  water:  1  part  in  600 
parts  of  water  at  18°,  about  twice  as  much  water  being  required  at 
100°.  The  solution,  relatively  to  its  concentration,  is  strongly 
alkaline.  On  account  of  its  cheapness,  this  substance  is  used  by 


478  COLLEGE    CHEMISTRY 

manufacturers  in  almost  all  operations  requiring  a  base,  and  it 
thus  occupies  the  same  position  amongst  bases  that  sulphuric  acid 
does  amongst  acids.  Caustic  lime  is  employed  in  the  manu- 
facture of  alkalies  (p.  446),  bleaching  powder,  and  mortar  (see 
below),  the  removal  of  the  hair  from  hides  in  preparation  for 
tanning,  the  softening  of  water  (see  below)  and  as  a  whitewash. 

Mortar.  —  Mortar  is  made  by  mixing  water  with  slaked  lime 
and  a  large  proportion  of  sand.  The  " hardening"  process  con- 
sists in  an  interaction  of  the  carbon  dioxide  of  the  air  with  the 
calcium  hydroxide: 

C02  +  Ca(OH)2  -»  CaC03  +  H20. 

After  the  superficial  parts  have  been  changed,  the  process  goes  on 
very  slowly,  and  many  years  are  required  before  the  deeper 
layers  have  been  transformed.  The  minute  crystals  of  calcite 
which  are  formed  are  interlaced  with  the  sand  particles,  and  a 
rigid,  yet  porous  mass  is  produced.  The  "hardening"  does  not 
begin  until  the  excess  of  water  used  in  making  the  mortar  has 
evaporated,  and  hence  ordinary  mortar  is  unsuitable  for  use  in 
damp  places  such  as  cellars. 

Calcium  Oxalate  CaC2O4.  —  This  salt  may  be  observed  under 
the  microscope  in  the  cells  of  many  plants.  It  appears  in  the 
form  of  needle-shaped  or  of  granular  crystals.  Since  it  is  the 
least  soluble  salt  of  calcium,  its  formation  by  precipitation  is  used 
as  a  test  for  calcium  ions. 

Theory  of  Precipitation.  —  The  precipitation  of  calcium  oxa- 
late  CaC204,  just  referred  to,  is  a  typical  one  and  may  be  used  to 
illustrate  the  application  of  ion-product  constancy  (p.  470)  to 
explaining  the  phenpmenon.  The  same  explanation  serves  for  all 
precipitations  of  ionogens. 

The  first  thing  to  be  remembered  is  that  the  precipitate  which 
we  observe,  however  insoluble  its  material  may  be,  does  not 
include  all  of  the  substance,  but  only  the  excess  beyond  what  is 
required  to  saturate  the  water.  The  liquid  surrounding  the  pre- 
cipitate is  always  a  saturated  solution  of  the  substance  precipitated. 
If  it  were  not  so,  some  of  the  precipitate  would  dissolve  until  the 


THEORY   OF   PRECIPITATION  479 

liquid  became  saturated.    Thus,  for  example,  when  we  add  am- 
monium oxalate  solution  to  calcium  chloride  solution:. 


(NH4)2C204  fc*  2NH4+- +  C£>*=  j  ^  CaC2o4  (dslvd)^  CaQA  (solid), 
CaCl2  £*  2C1    T  ^&          ) 

the  liquid  is  a  saturated  solution  of  calcium  oxalate,  with  the  excess 
of  this  salt  suspended  in  it  as  a  precipitate. 

Looking  at  the  matter  from  this  view  point,  we  perceive  the  ap- 
plication of  the  rule  of  ion-product  constancy.  In  this  saturated 
solution  (p.  470)  the  product  of  the  ion-concentrations,  [Ca++]  X 
[C204=],  is  constant.  If  the  original  solutions  had  been  so  very 
dilute  that,  when  they  were  mixed,  the  product  of  the  concentra- 
tions of  these  two  ions  had  not  reached  the  value  of  this  constant, 
no  precipitation  would  have  occurred.  As  a  matter  of  fact  the  ion- 
product  considerably  exceeded  the  requisite  value,  and  hence  the 
salt  was  thrown  down  until  the  balance  remaining  gave  the  value 
in  question.  The  rule  for  precipitation,  then,  is  as  follows:  When- 
ever the  product  of  the  concentrations  of  any  two  ions  in  a  mixture 
exceeds  the  value  of  the  ion-product  in  a  saturated  solution  of  the 
compound  formed  by  their  union,  this  compound  will  be  precipi- 
tated. Naturally  the  substances  with  small  solubilities,  and  there- 
fore small  ion-product  constants,  are  the  ones  most  frequently 
formed  as  precipitates. 

In  the  case  of  calcium  oxalate,  the  molar  solubility  (see  Table) 
is  0.0443.  In  so  dilute  a  solution  the  substance,  being  a  salt 
(p.  242),  must  be  practically  all  ionized.  Each  molecule  gives 
one  ion  of  each  kind.  The  molar  concentrations  of  these  ionic 
substances,  Ca++  and  C204=,  in  the  solution,  when  the  solid  is 
also  present,  must  therefore  be  (practically)  0.0443,  each.  The 
product  [Ca++]  X  [C204=]  is  thus  equal  to  0.0443  X  0.0443  or 
0.08185.  If  in  mixing  the  solutions,  exactly  equivalent  quan- 
tities were  not  employed,  the  values  of  the  two  factors  will  not 
be  equal,  but  the  product  will  in  any  case  possess  this  value. 

Rule  for  Solution  of  Substances.  —  The  rule  for  solution  of 

any  ionre  n  follows  at  once  from  the  foregoing  considerations,  and 
may  be  iormulated  by  changing  a  few  of  the  words  in  the  rule  just 
given:  Whenever  the  product  -f  the  concentrations  of  any  two 
ions  in  a  mixture  is  less  than  x  e  vai "S  "  the  ion-product  in  a 


480  COLLEGE    CHEMISTRY 

saturated  solution  of  the  compound  formed  by  their  union,  this 
compound,  if  present  in  the  solid  form,  will  be  dissolved.  When 
applied  to  the  simplest  case,  this  rule  means  that  a  substance  will 
dissolve  in  a  liquid  not  yet  saturated  with  it,  but  will  not  dissolve 
in  a  liquid  already  saturated  with  the  same  material.  The  value 
of  the  rule  lies  in  its  application  to  the  less  simple,  but  equally 
common  cases,  such  as  when  an  insoluble  body  is  dissolved  by 
interaction  with  another  substance  (next  section). 

Applications  of  the  Rule  for  Solution  to  the  Solution  of 
Insoluble  Substances.  —  So  long  as  a  substance  remains  in  pure 
water  its  solubility  is  fixed.  Thus,  with  calcium  hydroxide,  the 
system  comes  to  equilibrium  at  18°  when  0.17  g.  per  100  c.c.  of 
water  (0.02  moles  per  liter)  have  gone  into  solution : 

Ca(OH)2  (solid)  <=>  Ca(OH)2  (dslvd)  ±=>  Ca++  +  20H~. 

But  if  an  additional  reagent  which  can  combine  with  either  one  of 
the  ions  is  added,  the  concentration  of  this  ion  at  once  becomes  less, 
the  actual  numerical  value  of  the  ion-product  therefore  begins  to 
diminish,  and  further  solution  must  take  place  to  restore  its 
value.  Thus,  if  a  little  of  an  acid  (giving  H+)  be  added  to  the 
solution  of  calcium  hydroxide,  the  union  of  OH~  and  H+  to  form 
water  removes  the  OH~,  and  solution  of  the  hydroxide  proceeds 
until  the  acid  is  used  up.  There  are  now  more  Ca++  than  OH~ 
ions  present,  but  the  ion-product  reaches  the  same  value  as  be- 
fore, and  then  the  change  ceases.  If  a  further  supply  of  acid  is 
added,  the  removal  of  OH~  to  form  H20  begins  again.  With 
excess  of  the  acid,  the  only  stable  OH~  concentration  is  that  which 
is  a  factor  in  the  very  minute  ion-product  of  water,  [OH~]  X  [H+], 
which  is  0.061  X  0.061,  or  0.0i3l.  Hence,  with  excess  of  acid,  the 
calcium  hydroxide,  which  requires  in  general  a  much  higher  con- 
centration of  OH  than  this  to  precipitate  it  or  to  keep  it  out  of 
solution,  finally  all  dissolves. 

More  specifically,  if  we  assume  that  the  calcium  hydroxide  is 
wholly  dissociated  in  so  dilute  a  solution  (which  is  nearly  true), 
each  molecule  forms  one  ion  of  Ca++  and  two  ions  of  OH~~.  That 
is,  each  mole  of  Ca(OH)2  gives  one  mole  of  Ca++  and  two  moles 
of  OH  .  As  the  saturated  solution  contains  0.02  moles  of  the 
base,  the  molar  concentration  (assuming  complete  dissociation) 


THEORY   OF   PRECIPITATION 


481 


of  [Ca++]  is  0.02  and  of  [OHT]  is  0.04.  Now,  the  ion-product  is 
the  product  of  the  concentrations  of  all  the  ions  formed,  i.e. 
Ca++  OH~,  and  OH~.  The  value  of  the  product  is  therefore 
[Ca++]  X  [OH"]  X  [OH"]  or  [Ca++]  X  [OH~]2.  That  is,  0.02  X 
0.042  =  0.0432.  Note  that  if  the  molecule  gives  two  (or  three)  ions 
of  the  same  kind,  the  whole  concentration  of  that  ion  is  taken,  and 
is  also  raised  to  the  second  (or  third)  power. 

This  particular  action  is  a  neutralization  of  an  insoluble  base. 
But  the  other  kinds  of  actions  by  which  insoluble  ionogens  pass 
into  solution  all  resemble  it  closely,  and  differ  only  in  details.  The 
general  outlines  of  the  explanation  are  the  same  in  every  case. 
We  proceed  now  to  apply  it  to  the  common  phenomenon  of  the 
solution  of  an  insoluble  salt  by  an  acid. 

Interaction  of  Insoluble  Salts  with  Acids,  Resulting  in 
Solution  of  the  Salt.  —  Calcium  oxalate  passes  into  solution 
when  in  contact  with  acids,  especially  active  acids.  ^  Thus,  with 
hydrochloric  acid,  it  gives  calcium  chloride  and  oxalic  acid,  both 
of  which  are  soluble: 

CaC204  T  +  2HC1  fc?  CaCl2  +  H2C2O4.  (1) 

The  action  of  acids  upon  insoluble  salts  is  so  frequently  mentioned 
in  chemistry  and  is  so  important  a  factor  in  analytical  operations 
that  it  demands  separate  discussion.  This  example  is  a  typical 
one  and  may  be  used  as  an  illustration. 

According  to  the  rules  already  explained  (p.  479),  calcium 
oxalate  (or  any  other  salt)  is  precipitated  when  the  numerical  value 
of  the  product  of  the  concentrations  of  the  two  requisite  ions 
[Ca++]  X  [C204=]  exceeds  the  value  of  the  ion-product  ^for  a 
saturated  solution  of  calcium  oxalate  in  pure  water,  that  is,  ex- 
ceeds 0.08185  (p.  479).  When,  on  the  contrary,  the  product  of 
the  concentrations  of  the  two  ions  falls  below  the  limiting  value, 
a  condition  which  may  arise  from  the  removal  in  some  way  either 
of  the  Ca++  or  of  the  C204=  ions,  the  undissociated  molecules  will 
ionize,  and  the  solid  will  dissolve  to  replace  them  until  the  ionic 
concentrations  necessary  for  equilibrium  with  the  molecules  have 
been  restored  or  until  the  whole  of  the  solid  present  is  consumed. 
Here  the  oxalate-ion  from  the  calcium  oxalate  combines  with  the 


482  COLLEGE    CHEMISTRY 

hydrogen-ion  of  the  acid  (usually  an  active  one)  which  has  been 
added,  and  forms  molecular  oxalic  acid: 

(2) 


Hence,  dissociation  of  the  dissolved  molecules  of  calcium  oxalate 
proceeds,  being  no  longer  balanced  by  encounters  and  unions  of 
the  now  depleted  ions,  and  this  dissociation  in  turn  leads  to  solu- 
tion of  other  molecules  -from  the  precipitate. 

It  will  be  seen  that  the  removal  of  the  ions  in  this  fashion  can 
result  in  considerable  solution  of  the  salt  only  when  the  acid  pro- 
duced is  a  feebly  ionized  one.  Here,  to  be  specific,  the  concentra- 
tion of  the  C204=  in  the  oxalic  acid  equilibrium,  (2)  above,  must  be 
less  than  that  of  the  same  ion  in  a  saturated  calcium  oxalate 
solution.  Now  oxalic  acid  does  not  belong  to  the  least  active 
class  of  acids,  and  its  pure  solution  contains  a  considerable  con- 
centration of  C2O4=.  There  is,  however,  a  decisive  factor  in  the 
situation  which  we  have  not  yet  taken  into  account.  The  hydro- 
chloric acid  which  we  used  for  dissolving  the  precipitate  is  a  very 
highly  ionized  acid  and  gives  an  enormously  greater  concentration 
of  hydrogen-ion  than  does  oxalic  acid.  Hence  the  hydrogen-ion 
is  in  excess  in  equation  (2),  and  the  condition  for  equilibrium, 

[H+]2  X  [0,04=]       ^      .„  ,        ,  .  -    ,  , 

be  satisfied  by  a  correspondingly  small 


concentration  of  C2O4=.  In  this  particular  case,  therefore,  the 
[C204~]  of  the  oxalic  acid  is  less  than  that  given  by  the  calcium 
oxalate.  The  whole  change,  therefore,  depends  for  its  accomplish- 
ment, not  only  on  the  mere  presence  of  hydrogen-ion,  but  on  the 
repression  of  the  ionization  of  the  oxalic  acid  by  the  great  excess  of 
hydrogen-ion  furnished  by  the  active  acid  that  has  been  used.  As 
a  matter  of  fact,  we  find  that  a  weak  acid  like  acetic  acid  has 
scarcely  any  effect  upon  a  precipitate  of  calcium  oxalate.  An 
acid  stronger  than  oxalic  acid  must  be  employed.  The  whole 
scheme  of  the  equilibria  is  as  follows: 


2HC1 

When  excess  of  an  acid  sufficiently  active  to  furnish  a  large  con- 
centration of  hydrogen-ion  is  employed,  the  last  equilibrium  is 
then  driven  forward  and  the  others  follow.  With  addition  of  a 


THEORY   OF    PRECIPITATION  483 

weak  acid,  only  a  slight  displacement  occurs,  and  the  system  comes 
to  rest  again  when  the  molecular  oxalic  acid  has  reached  a  sufficient 
concentration. 

A  generalization  may  be  based  on  these  considerations  :  an  insoluble 
salt  of  a  given  acid  will  in  general  interact  and  dissolve  when  treated 
with  a  solution  containing  another  acid,  provided  that  the  latter  acid 
is  a  much  more  highly  ionized  (more  active)  one  than  the  former 
(see  below)  . 

But  even  active  acids  frequently  fail  to  bring  salts  of  weak  acids 
into  solution,  especially  when  the  weak  acid  is  itself  present  also. 
Here  the  cause  lies  in  the  fact  that  such  salts  are  even  less  soluble 
than  those  of  the  calcium  oxalate  type,  and  give  so  low  a  con- 
centration of  the  negative  ion  that  the  utmost  repression  of  the 
ionization  of  the  corresponding  acid  does  not  give  a  lower  value 
for  the  concentration  of  this  ion  than  does  the  salt  itself.  Thus, 
we  have  seen  (p.  272)  that  even  hydrochloric  acid  (dilute)  will  not 
dissolve  a  number  of  sulphides.  For  example,  in  the  case  of 
cupric  sulphide  in  a  solution  saturated  with  hydrogen  sulphide, 
the  S=  factor  in  the  solubility  product  [Cu++]  X  [S=]  remains 
smaller  than  that  in  the  scheme  defining  the  hydrogen  sulphide 
equilibrium  [H+]2  X  [S=]  even  when  the  [S=]  factor  in  the  latter 
is  diminished  in  consequence  of  great  addition  of  hydrogen-ion. 
In  this  case  the  first  link  in  the  chain  of  equilibria: 

CuS  (solid)  <=»  CuS  (dslvd)  ±9  Cu++  +  S=  )  _>  TT  «  ,A  i    ~ 
2HC1  fc*2Cr  +  2H+r  slvd)> 


tends  so  decidedly  backward  that  only  the  use  of  concentrated  acid 
will  increase  the  concentration  of  the  H+  to  an  extent  sufficient  to 
secure  even  a  slight  advance  of  the  whole  action.  We  must  add, 
therefore,  to  the  above  rule  :  provided  also  that  the  salt  is  not  one 
of  extreme  insolubility.  This  point  will  be  illustrated  more  fully  in 
connection  with  the  description  of  individual  sulphides  (see  under 
Cadmium). 

Illustrations  of  the  application  of  these  generalizations  are 
countless.  Carbonic  acid  is  made  from  marble  (p.  381),  hydrogen 
sulphide  from  ferrous  sulphide  (p.  272),  hydrogen  peroxide  from 
sodium  peroxide  (p.  222),  and  phosphoric  acid  from  calcium  phos- 
phate (p.  370).  In  each  case  the  acid  employed  to  decompose  the 
salt  is  more  active  than  the  acid  to  be  liberated.  On  the  other 


484  COLLEGE   CHEMISTRY 

hand,  calcium  oxalate  is  insoluble  in  acetic  acid  because  this  acid 
is  weaker  than  is  oxalic  acid.  We  have  thus  only  to  examine  the 
list  of  acids  showing  their  degrees  of  ionization  (p.  241)  in  order 
to  be  able  to  tell  which  salts,  if  insoluble  in  water,  will  be  dis- 
solved by  acids  and,  in  general,  what  acids  will  be  sufficiently 
active  in  each  case  for  the  purpose.  In  chemical  analysis  we  dis- 
criminate between  salts  soluble  in  water,  those  soluble  in  acetic 
acid  (the  insoluble  carbonates  and  some  sulphides,  FeS  and  MnS, 
for  example),  those  requiring  active  mineral  acids  for  their  solu- 
tion (calcium  oxalate  and  the  more  insoluble  sulphides,  for  ex- 
ample), and  those  insoluble  in  all  acids  (barium  sulphate  and 
other  insoluble  salts  of  active  acids). 

Precipitation  of  Insoluble  Salts  in  Presence  of  Acids.  — • 

The  converse  of  solution,  namely,  precipitation,  depends  upon  the 
same  conditions:  an  insoluble  salt  which  is  dissolved  by  a  given  acid 
cannot  be  formed  by  precipitation  in  the  presence  of  this  acid.  Thus, 
calcium  oxalate  can  be  precipitated  in  presence  of  acetic  acid,  but 
not  in  presence  of  active  mineral  acids  in  ordinary  concentrations. 
Cupric  sulphide  or  barium  sulphate  can  be  precipitated  in  pres- 
ence of  any  acid,  but  ferrous  sulphide  and  calcium  carbonate  only 
in  the  absence  of  acids. 

From  this  it  does  not  follow  that  calcium  oxalate,  for  example, 
cannot  be  precipitated  if  once  an  active  acid  has  been  added  to  the 
mixture.  To  secure  precipitation,  all  that  is  necessary  is  to  re- 
move the  excess  of  hydrogen-ion  which  is  repressing  the  ionization 
of  the  oxalic  acid.  This  can  be  done  by  adding  a  base,  which  re- 
moves the  H+,  or  even  by  adding  sodium  acetate.  The  acetate- 
ion  C2H302~  unites  with  the  H+  to  form  the  little  ionized  acetic 
acid,  in  presence  of  which  calcium  oxalate  can  be  precipitated. 

Bleaching  Powder  Ca(OCl)CL  —  This  substance  (cf.  p.  309) 
is  manufactured  by  conducting  chlorine  into  a  box-like  structure 
containing  slaked  lime  spread  upon  perforated  shelves.  When  the 
transformation  has  reached  the  limit  (it  is  never  complete),  some 
lime  dust  is  blown  into  the  chamber  to  absorb  the  remainder  of 
the  free  chlorine. 

That  bleaching  powder  is  a  mixed  salt  CaCl(OCl)  rather  than  an 
equimolar  mixture  of  calcium  chloride  and  calcium  hypochlorite, 


CALCIUM  485 

which  would  have  the  same  composition,  CaCl2,Ca(OCl)2,  is 
proved  by  the  facts  that  the  material  is  not  deliquescent  as  is 
calcium  chloride,  and  that  calcium  chloride  cannot  be  dissolved  out 
of  it  by  alcohol. 

Bleaching  powder  is  somewhat  soluble  in  water,  and  in  solution 
the  ions  Ca++,  Cl~,  and  C10~  are  all  present.  Addition  of  active 
acids  causes  the  formation  of  hydrochloric  and  hypochlorous  acids 
(p.  309).  Weak  acids  like  carbonic  acid  displace  the  hypo- 
chlorous  acid  only  (cf.  p.  310),  and  hence  the  dry  powder,  when 
exposed  to  the  air,  has  the  odor  of  hypochlorous  anhydride  C12O 
rather  than  that  of  chlorine. 

The  substance  is  largely  used  by  bleachers  (cf.  p.  311),  and  as  a 
disinfectant  to  destroy  germs  of  putrefaction  and  disease. 

Calcium  Sulphate.  —  This  salt  is  found  in  large  quantities  in 
nature.  The  mineral  anhydrite  CaSO4  occurs  in  the  salt  layers. 
It  contains  no  water  of  hydration,  and  its  crystals  belong  to  the 
rhombic  system.  The  dihydrate,  CaSO4,2H2O,  is  more  plentiful. 
In  granular  masses  it  constitutes  alabaster.  When  perfectly  crys- 
tallized (monoclinic  system,  Fig.  47,  p.  83)  it  is  named  gypsum  or 
selenite.  The  same  hydrate  is  formed  by  precipitation  from  solu- 
tions. Its  solubility  is  about  1  in  500  at  18°. 

Plaster  of  paris  2CaS04,H20  is  manufactured  by  heating  gyp- 
sum until  nearly  all  the  water  of  hydration  has  been  driven 
out.  When  it  is  mixed  with  water,  the  dihydrate  is  quickly  re- 
formed and  a  rigid  mass  is  produced.  That  the  plaster  sets 
rapidly,  is  due  to  the  fact  that  the  hemihydrate  is  more  soluble 
than  the  dihydrate,  and  so  a  constant  solution  of  -the  one  and 
deposition  of  the  other  goes  on  until  the  hydration  is  complete. 
It  becomes  rigid,  instead  of  forming  a  loose  mass  of  dihydrate, 
because  the  process  results  in  the  formation  of  an  interlaced  and 
coherent  mass  of  minute  crystals. 

2CaS04(H20  (solid)  ^2CaSO4  (dslvd)  j  ^  2CaSO4)2H2O  (solid). 

Plaster  of  paris  is  used  for  making  casts  and  in  surgery.  The 
setting  of  the  material  is  accompanied  by  a  slight  increase  in 
volume,  and  hence  a  very  sharp  reproduction  of  all  the  details  of 
the  mold  is  obtained.  An  "ivory"  surface,  which  makes  washing 


486  COLLEGE    CHEMISTRY 

practicable,  is  conferred  by  painting  the  cast  with  a  solution  of 
paraffin  or  stearin  in  petroleum  ether.  The  waxy  material,  left 
by  evaporation  of  the  volatile  hydrocarbons,  fills  the  pores  and 
prevents  solution  and  disintegration  of  the  substance  by  water. 
Stucco  is  made  with  plaster  of  paris  and  rubble,  and  is  mixed  with 
a  solution  of  glue  instead  of  water. 

Calcium  Sulphide  CaS.  —  This  compound  is  most  easily  made 
by  strongly  heating  pulverized  calcium  sulphate  and  charcoal. 
The  sulphate  is  reduced:  4C  +  CaSO4  ->  CaS  +  4CO.  Calcium 
sulphide  is  meagerly  soluble  in  water,  but  is  nevertheless  slowly 
dissolved  in  consequence  of  its  decomposition  by  hydrolysis 
into  calcium  hydroxide  and  calcium  hydrosulphide  Ca(SH)2. 
It  is  used  as  a  depilatory.  Hair  and  wool  are  composed  of  proteins, 
which  are  decomposed  by,  and  dissolved  in  alkaline  solutions, 
like  that  here  formed.  Since  calcium  sulphide  is  thus  decomposed 
by  water  it  cannot  be  precipitated  from  aqueous  solution  by 
adding  a  soluble  sulphide. 

Ordinary  calcium  sulphide,  after  it  has  been  exposed  to  sunlight, 
usually  shines  in  the  dark.  Barium  sulphide  behaves  in  the  same 
way.  On  this  account  these  substances  are  used  in  making 
luminous  paint.  They  apparently  owe  this  behavior  to  the 
presence  of  traces  of  compounds  of  vanadium  and  bismuth,  for 
the  purified  substances  are  not  affected  in  the  same  fashion. 

Phosphates  of  Calcium.  —  The  tertiary  orthophosphate  of 
calcium  Ca3(P04)2,  known  as  phosphorite,  is  found  in  many  locali- 
ties, and  is  often  derived  from  the  remains  of  animals.  Guano 
contains  some  of  the  same  substance,  along  with  nitrogen  either 
in  the  form  of  organic  compounds  or  as  niter  (p.  348).  Apatite, 
3Ca3(PO4)2,CaF2,  a  double  salt  with  calcium  fluoride  (or  chloride), 
is  a  common  mineral  and  frequent  component  of  rocks.  The 
orthophosphate  forms  about  83  per  cent  of  bone  ash,  "and  is 
contained  also  in  the  ashes  of  plants.  It  may  be  precipitated 
by  adding  a  soluble  phosphate  to  a  solution  of  a  salt  of 
calcium. 

Since  it  is  a  salt  of  a  weak  acid,  and  belongs  to  the  less  insoluble 
class  of  such  salts,  calcium  phosphate  is  dissolved  by  dilute  mineral 
acids  (cf.  p.  483),  the  ions  HPO4=  and  H2PO4~  being  formed. 


CALCIUM  487 

When  a  base,  such  as  ammonium  hydroxide,  is  added  to  the  solu- 
tion, the  calcium  phosphate  is  reprecipitated  (cf.  p.  484) . 

Calcium  phosphate  is  chiefly  used  in  the  manufacture  of  phos- 
phorus and  phosphoric  acid  (p.  370),  and  as  a  fertilizer.  The 
supply  of  calcium  phosphate  in  the  soil  arises  from  the  decompo- 
sition of  rocks  containing  phosphates,  and  is  gradually  exhausted 
by  the  removal  of  crops.  Bone  ash  is  sometimes  used  to  make 
up  the  deficiency.  It  is  almost  insoluble  in  water,  however,  and, 
although  somewhat  less  insoluble  in  natural  water  containing 
salts  like  sodium  chloride,  is  brought  into  a  condition  for  absorp- 
tion by  the  plants  rather  slowly.  The  "superphosphate"  (see 
below)  is  much  more  soluble. 

Primary  calcium  orthophosphate  ("superphosphate")  is  manu- 
factured in  large  quantities  from  phosphate  rock  by  the  action  of 
sulphuric  acid.  The  unconcentrated  "chamber  acid"  is  used  for 
this  purpose,  as  water  is  required  in  the  resulting  action.  The 
amounts  of  material  employed  correspond  to  the  equation: 

Ca3(P04)2  +  2H2S04  +  6H20  -*  Ca(H2P04)2,2H2O  +  2CaSO4,2H20. 

As  soon  as  mixture  has  been  effected,  the  action  proceeds  with 
evolution  of  heat,  and  a  large  cake  of  the  two  hydrated  salts  re- 
mains. This  mixture,  after  being  broken  up,  dried,  and  packed 
in  bags,  is  sold  as  "superphosphate  of  lime."  The  primary  phos- 
phate which  it  contains  is  soluble  in  water,  and  is  therefore  of 
great  value  as  a  fertilizer. 

Calcium  Cyanamide.  —  Calcium  carbide,  (p.  379),  when 
strongly  heated  with  nitrogen,  gives  a  mixture  of  calcium  cyan- 
amide  and  carbon: 

CaC2  +  N2-»CaCN2  +  C, 

which  is  sold  as  nitre-lime  for  use  as  a  fertilizer.  When  treated 
with  hot  water,  the  cyanamide  is  hydrolyzed  into  calcium  carbon- 
ate and  ammonia: 

CaCN2  +  3H20  -» CaC03  +  2NH3. 

In  the  soil  the  decomposition  may  not  be  so  simple,  but  combined 
nitrogen  is  furnished  in  a  form  that  can  be  absorbed  by  plants. 


488 


COLLEGE    CHEMISTRY 


At  Odda  (Norway)  the  carbide  is  pulverized  and  placed  in  a 
cylindrical  furnace  (Fig.  119)  holding  300-450  kg.     The  heat  (800- 
1000°)  is  supplied  by  the  passage  of  an  electric  current  through 
a  thin  carbon  rod.     The  nitrogen"  is  obtained  by 
the  fractionation  of  liquid  air  and  final  removal 
of  all  oxygen  by  passage  over  heated  copper,  and 
is  forced  in  under  pressure.     After  thirty-six  hours, 
nitrogen  is  no  longer  absorbed,  and  the  charge  is 
pulverized  when  cold. 

Sodium  cyanide  is  now  manufactured  by  fusing 
nitro-lime  with  sodium  carbonate: 


CaCN2  +  C 


CaCO3  +  2NaNC. 


FIG.  119. 


The  cyanide  is  extracted  from  the  insoluble  cal- 
cium carbonate  with  water,  in  which  it  is  exceed- 

ingly soluble.   Sodium  cyanide  has  now  displaced  potassium  cyanide 

in  the  extraction  of  gold  from  its  ores. 


Nutrition  and  Fertilization  of  Crops.  —  A  plant  constructs 
its  cellulose,  starch,  and  sugar,  and  secures  the  carbon-part  of 
all  its  organic  contents  from  the  carbon  dioxide  of  the  air  (p.  387). 
The  water  (90-95  per  cent  of  the  total  weight  of  the  plant)  comes 
from  the  soil  and  brings  up  in  solution  the  other  elements  re- 
quired. All  soils  are  able  to  supply  sufficient  magnesium,  calcium 
and  iron  as  bicarbonates.  But  the  soil  may  lack:  sulphur, 
absorbed  as  sulphates;  nitrogen,  absorbed  chiefly  as  nitrates,  but 
occasionally  as  salts  of  ammonium;  potassium,  as  sulphate, 
chloride,  or  nitrate;  and  phosphorus,  as  soluble  phosphates.  The 
soil  may  be  originally  deficient  in  one  or  more  of  these  necessary 
plant  foods,  or  the  supply  may  have  been  exhausted  by  repeated 
cropping.  Every  crop  permanently  removes  certain  quantities. 
For  example,  in  the  case  of  nitrogen,  which  is  required  to  form 
proteins  that  enter  largely  into  the  fruit  (i.e.,  usually,  the  edible 
part),  each  crop  of  Indian  corn  (45  bushels)  removes  63  pounds 
per  acre,  a  crop  of  cabbage  (15  tons)  removes  100  pounds  per  acre, 
clover  hay  (2  tons)  82  pounds,  and  wheat  (15  bushels)  31  pounds. 
When  the  store  in  the  soil  become  meager,  the  crops  become  poor, 
and  finally  cost  more  for  labor  than  they  are  worth. 


CALCIUM  489 

Thus,  crops  have  to  be  fed,  just  like  cattle.  Moreover,  the 
elements  must  be  furnished  in  soluble  form  (cf.  pp.  137,  340,  422). 
Fertilizers  containing  potassium  (p.  445,  451)  and  phosphorus 
(p.  487)  must  be  used,  when  the  soil  is  deficient  in  these  elements. 
The  nitrogen  fertilizers  we  have  mentioned  are  sodium  nitrate 
(p.  347),  calcium  nitrate  (p.  353),  ammonium  sulphate  (p.  411), 
guano  and  manure  (p.  348),  " tankage"  and  ground  bones  from 
slaughter  houses,  calcium  cyanamide  (p.  487),  and  finally  the 
nitrates  from  bacterial  decomposition  of  root  nodules  (p.  339). 
That  systematic  use  of  fertilizers  does  influence  the  crops  is  indi- 
cated by  the  results  of  cultivation  of  land  which,  but  for  fertili- 
zation, would  long  since  have  become  almost  valueless.  The 
wheat  crop  per  acre,  being  the  average  of  ten  successive  years  is: 
Denmark  40  bushels,  Great  Britain  33,  Germany  29,  United 
States  14. 

Hard  Water.  —  As  we  have  seen  (pp.  384,  476),  limestone 
(solubility,  0.013  g.  per  liter),  magnesium  carbonate  (sol'ty  1  g. 
per  liter),  and  iron  carbonate,  although  very  insoluble,  are  acted 
upon  by  the  carbonic  acid  in  natural  waters,  giving  bicarbonates 
which  are  roughly  about  thirty  times  as  soluble.  When  the 
water  is  boiled,  the  actions  are  reversed,  and  the  carbonates 
are  reprecipitated.  These  bicarbonates  constitute  temporary 
hardness,  and  their  decomposition  produces  "fur"  in  a  kettle  and 
boiler  crust  in  a  boiler. 

The  sulphates  of  calcium  (sol'ty  2  g.  per  liter)  and  of  magnesium 
(sol'ty  354  g.  per  1)  are  also  commonly  found  in  natural  waters. 
These  salts  are  not  affected  by  mere  boiling  (as  distinct  from 
evaporation)  and  so,  along  with  magnesium  carbonate  (1  g.  per  1.) 
and  calcium  carbonate  (0.013  g.  per  1.)  give  permanent  hardness 
to  the  water. 

Hardness  is  estimated  in  "degrees."  In  France,  and  com- 
monly in  the  laboratory,  1  part  of  CaC03  (or  its  equivalent  of 
other  salts)  per  100,000  (0.01  g.  per  liter)  constitutes  one  degree. 
In  the  United  States  one  degree  is  1  grain  per  gallon  of  58,333 
grains  (0.017  g.  per  1.).  In  Britain  one  degree  is  1  grain  per  gallon 
of  70,000  grains  (0.014  g.  per  1.).  Well  water,  originating  in 
chalk  or  limestone  formations,  may  have  37°  (Fr.)  or  more  of 
hardness. 


490  COLLEGE   CHEMISTRY 

Damage  Due  to  Hardness  in  Water.  —  When  hard  water  is 
continually  fed  into  a  steam  boiler  and  only  steam  comes  out, 
naturally  the  salts  accumulate  and  produce  in  time  a  heavy 
boiler  crust,  which  settles  on  the  tubes.  Being  a  poor  conductor 
of  heat  compared  with  iron,  this  crust,  if  one-fourth  of  an  inch 
thick,  will  increase  the  consumption  (and  cost)  of  fuel  by  50  per 
cent.  In  addition,  the  iron,  not  being  in  direct  contact  with 
water,  is  heated  to  a  higher  temperature,  and  may  even  become 
red  hot.  It  thus  oxidizes  more  quickly  on  the  outside,  and  dis- 
places hydrogen  from  water  (or  steam)  on  the  inside  (p.  52), 
thus  changing  on  both  sides  gradually  into  the  brittle  magnetic 
oxide  Fe3O4.  If  the  crust  is  not  removed,  or  prevented  (see 
below),  the  life  of  the  boiler  is  greatly  shortened,  and  a  serious 
explosion  may  even  occur. 

In  washing,  in  the  household  or  laundry,  much  soap  is  wasted 
before  the  necessary  lather  is  secured.  The  soap,  for  example, 
the  sodium  stearate  (p.  415),  gives  magnesium  and  calcium  stear- 
ates,  which  are  insoluble,  forming  a  curd: 

CaS04  +  2Na(C02C17H35)  -»  Ca(C02C17H35)2  1  +  Na2S04. 


The  permanent  solution  of  soap,  required  for  washing,  does  not 
begin  to  be  formed  until  all  the  hardness  has  thus  been  precipi- 
tated. Hence,  according  to  the  equation,  with  1°  (U.  S.)  hard- 
ness, 100  gallons  (U.  S.)  of  water  should  use  up  0.075  pound  of 
soap  (1°  Brit,  and  100  gal.  Brit.,  0.075  lb.).  In  point  of  fact, 
however,  the  colloidal  calcium  salts  adsorb  and  carry  down  with 
them  more  than  an  equal  amount  of  undecomposed  soap.  Hence, 
actual  measurement  shows  that,  with  1°  (U.  S.  or  Brit.)  of  hard- 
ness, 100  gallons  (U.  S.  or  Brit.)  of  water  really  destroy  0.17 
pound  of  soap.  Thus,  with  35°,  no  less  than  6  pounds  of  soap 
per  100  gallons  are  wasted  before  the  part  of  the  soap  that  is  to 
do  the  work  begins  to  dissolve. 

Treatment  of  Hard  Water.  —  The  temporary  hardness  can 
be  removed  by  boiling  the  water,  or  using  some  preheating 
arrangement  in  connection  with  the  boiler  (stationary  engines 
only). 

Temporary  hardness  is  commonly  removed,  on  a  large  scale,  by 
adding  slaked  lime  (made  into  milk  of  lime)  in  exactly  the  quantity 


CALCIUM  491 

shown  by  an  analysis  of  the  water  to  be  required,  and  stirring  for 
a  considerable  time: 

Ca(HC03)2  +  Ca(OH)2  ->  2CaC03  J,  +  2H20.  (1) 

The  bicarbonate  is  neutralized  and  all  the  lime  precipitated. 
The  latter  is  removed  by  filtration. 

Permanent  hardness  is  not  affected  by  slaked  lime,  but  is  pre- 
cipitated by  adding  sodium  carbonate  in  the  necessary  proportion: 
CaS04  +  Na2C03  ->  CaCO3  J,  +  Na2S04.  (2) 

When  both  kinds  of  hardness  are  present,  crude  caustic  soda 
(sodium  hydroxide)  may  be  employed.  It  neutralizes  the  bicar- 
bonate, precipitating  CaCO3: 

Ca(HC03)2  +  2NaOH  -» CaC03  J,  +  Na2C03  +  H2O.        (3) 

and  giving  sodium  carbonate.  The  latter  then  acts  as  in  equa- 
tion (2). 

Instead  of  this,  the  treatments  indicated  in  equations  (1)  and 
(2)  may  be  applied  in  combination  (Porter-Clark  process).* 

In  the  new  permutite  process  the  water  is  simply  filtered  through 
an  artificial  sodium  silico-aluminate  (permutite  NaP)  which  is 
supplied  in  the  form  of  a  coarse  sand.     The  calcium,  etc.,  in  the 
water  is  exchanged  for  sodium,  which  does  no  harm: 
Ca(HC03)2  +  2NaP  -~»  2NaHC03  +  CaP2. 

After  twelve  hours'  use,  the  permutite  is  covered  with  10  per 
cent  salt  solution,  and  allowed  to  remain  for  the  other  twelve 
hours  of  the  day,  when  it  is  ready  for  employment  once  more: 

2NaCl  +  CaP2  ->  CaCl2  +  2NaP. 

Only  salt,  which  is  inexpensive,  is  consumed,  and  calcium  chloride 
solution  is  thrown  away.  Permutite  removes  magnesium,  iron, 
manganese,  and  other  elements  in  the  same  way.  The  life  of  a 
charge  is  said  to  be  over  twenty  years. 

Hard  Water  in  the  Laundry.  —  As  we  have  seen  (p.  490), 
soap  will  soften  water,  but  the  calcium  and  magnesium  salts  of  the 
soap  acids,  which  are  precipitated,  are  sticky,  and  soil  the  goods 

*  So  far  as  the  hardness  is  due  to  magnesium  bicarbonate,  a  double  propor- 
tion of  lime  must  be  added  to  precipitate  the  magnesium  as  hydroxide  (soTty 
0.01  g.  per  1.),  because  the  carbonate  is  too  soluble  (1  g.  per  1.). 


492  COLLEGE    CHEMISTRY 

being  washed.  Other  substances  that  soften  water  not  only  give 
non-adhesive  precipitates,  but  are  also  much  cheaper,  and  an  at- 
tempt is  generally  made  to  utilize  them.  The  use  of  slaked  lime 
is  impracticable  on  a  small  scale. 

Washing  soda  Na2C03,10H20  is  added  to  precipitate  both  kinds 
of  hardness: 

Ca(HC03)2  +  NajCOi->CaCO8  +  2NaHCO3, 
CaSO4  +  Na2C03  -»  CaCO3  +  Na2SO4. 

The  small  amounts  of  salts  of  sodium  which  remain  in  the  water 
have  no  action  on  soap. 

Household  Ammonia  NH4OH  acts  like  sodium  hydroxide 
(p.  491): 

Ca(HC03)2  +  2NH4OH  -*  CaCO3  +  (NH^COs  +  2H20, 
CaSO4  +  (NH4)2C03-+CaC03  +  (NH4)2SO4. 

except  that  it  will  not  precipitate  magnesium-ion. 

Borax  Na2B406,10H20  (p.  432)  is  hydrolyzed  and  the  sodium 
hydroxide  contained  in  its  solution  acts  as  already  (p.  491)  de- 
scribed. 

The  supposed  bleaching  or  whitening  action  of  borax  or  soda 
is  a  myth;  these  salts  prevent  staining  by  the  iron  in  the  water. 
They  simply  precipitate  the  iron,  present  as  Fe(HCO3)2,  which 
almost  all  waters  contain,  as  FeCO3,  before  the  goods  are  put  in. 
This  precipitate  is  easily  washed  out  in  rinsing.  The  palmitate, 
etc.,  of  iron,  however,  which  the  soap  itself  would  throw  down,  is 
sticky  and  adheres  to  the  cloth.  The  air  subsequently  oxidizes 
it  (see  p.  633)  and  gives  hydrated  ferric  oxide  (rust),  which  is 
brownish-red. 

It  is  evident  that,  properly  to  achieve  their  purpose,  the  soda 
and  borax  must  be  added,  must  be  completely  dissolved,  and 
must  be  allowed  to  produce  the  precipitation  of  FeC03,  CaCO3, 
etc.,  all  before  the  soap  (or  the  goods)  is  introduced.  If  the  soap 
is  dissolved  before  or  with  the  soda,  it  will  take  part  in  the  pre- 
cipitation, and  give  sticky  particles  containing  the  iron  and  cal- 
cium salts  of  the  soap  acids. 

The  soda,  borax,  and  ammonia  do  not  themselves  remove  dirt 

-that  is  done  by  the  dissolved  soap  (p.  418).     With  the  help  of 

rubbing,  however,  they  do  emulsify  and  remove  animal  or  vege- 


CALCIUM  493 

table  oil  and  grease,  but  not  mineral  oil  (p.  420),  when  these 
happen  to  be  on  the  goods.  But  soap  alone  will  do  this  also,  and 
remove  mineral  oil  as  well. 

Washing  powders  are,  or  ought  to  be,  mainly  sodium  carbonate, 
mixed  with  more  or  less  pulverized  soap. 

Calcium  Silicate  CaSiO3.  —  Calcium  metasilicate  CaSi03 
forms  the  mineral  wollastonite,  which  is  rather  scarce,  but  enters 
into  the  composition  of  many  complex  minerals,  such  as  garnet 
and  mica.  It  may  be  made  by  precipitation  with  a  solution  of 
sodium  metasilicate  (p.  428),  or  by  fusing  together  powdered 
quartz  and  calcium  carbonate: 

Si02  +  CaCO3  -*  CaSi63  +  C02. 

Glass.  —  Common  glass  is  a  complex  silicate  of  sodium  and  cal- 
cium, or  a  homogeneous  mixture  of  the  silicates  of  these  metals 
with  silica.  It  has  a  composition  represented  approximately  by 
the  formula  Na2O,CaO,6Si02,  and  is  made  by  melting  together 
sodium  carbonate,  limestone,  and  pure  sand: 

Na2C03  +  CaC03  +  6Si02  -*  Na20,CaO,6Si02  +  2C02. 

For  the  most  fusible  glass,  a  smaller  proportion  of  sand  is  em- 
ployed. This  variety  is  known,  from  its  components,  as  soda- 
glass,  or,  from  its  easy  fusibility,  as  soft  glass.  Plate-glass  is  made 
by  casting  the  material  in  large  sheets,  rolling  the  sheet  flat  while 
hot,  and  polishing  the  surfaces  when  cold  until  they  are  plane. 
Window-glass  is  prepared  by  blowing  bulbs  of  long  cylindrical 
shape,  and  ripping  them  down  one  side  with  the  help  of  a  diamond. 
The  resulting  curved  sheets  are  then  placed-  on  a  flat  surface  in  a 
furnace  and  are  there  allowed  to  open  out.  Beads  are  made, 
chiefly  in  Venice,  by  cutting  narrow  tubes  into  very  short  sections 
and  rounding  the  sharp  edges  by  fire.  Ordinary  apparatus  is 
made  of  soft  soda-glass,  and  hence  when  heated  strongly  it  tends 
to  soften  and  also  to  confer  a  strong  yellow  tint  (cf.  p.  465)  on  the 
flame.  Bottles  are  made  with  impure  materials,  and  owe  their 
color  chiefly  to  the  silicate  of  iron  which  they  contain.  In  all 
cases  the  articles  are  annealed  by  being  passed  slowly  through  a 
special  furnace  in  which  their  temperature  is  lowered  very  grad- 


494  COLLEGE    CHEMISTRY 

ually.  Glass  which  has  been  suddenly  chilled  is  in  a  state  of 
tension  and  breaks  easily  when  handled. 

Soft  glass  is  partially  dissolved  by  water.  When  powdered 
glass  is  shaken  with  water,  sodium  silicate  dissolves  at  once, 
and  in  amount  sufficient  to  give  an  alkaline  reaction  with  phenol- 
phthale'in  (cf.  p.  258). 

Bohemian,  or  hard  glass,  is  much  more  difficult  to  fuse  than 
soda-glass,  and  is  also  much  less  soluble  in  water.  It  is  manu- 
factured by  substituting  potassium  carbonate  for  the  sodium 
carbonate.  Specially  insoluble  glass,  for  laboratory  use,  such  as 
Jena  and  non-sol  glass,  is  made  with  boric  anhydride  B2O3,  in  ad- 
dition to  silica,  and  some  zinc  oxide,  so  that  it  contains  borates  as 
well  as  silicates.  When  lead  oxide  is  employed  instead  of  lime- 
stone, a  soda-lead  glass  known  as  flint  glass  is  produced.  This 
has  a  high  specific  gravity,  and  a  great  refracting  power  for  light, 
and  is  employed  for  making  glass  ornaments.  By  the  use  of 
grinding  machinery,  cut  glass  is  made  from  it.  Engraving  on 
glass  is  done  with  the  sand  blast. 

Colored  glass  is  prepared  by  adding  small  amounts  of  various 
oxides  to  the  usual  materials.  The  oxides  combine  with  the 
silica,  and  produce  strongly  colored  silicates.  Thus,  cobalt  oxide 
gives  a  blue,  chromium  oxide  or  cupric  oxide  a  green,  and  uranium 
oxide  a  yellow  glass.  Cuprous  oxide,  with  a  reducing  agent,  and 
compounds  of  gold,  give  the  free  metals,  suspended  in  colloidal 
solution  (p.  416)  in  the  glass,  and  confer  a  deep-red  color  upon  it. 
Milk-glass  contains  finely  powdered  calcium  phosphate  in  sus- 
pension, and  white  enamels  are  made  by  adding  stannic  oxide. 

Glass  is  a  typical  amorphous  substance  (pp.  266,  393).  From 
the  facts  that  it  has  no  crystalline  structure,  and  that  it  softens 
gradually  when  warmed,  instead  of  showing  a  definite  melting- 
point,  it  is  regarded  as  a  supercooled  liquid  of  extreme  viscosity. 
Most  single  silicates  crystallize  easily,  and  have  definite  freezing- 
(and  melting-)  points.  Glass  may  be  regarded  as  a  solution  of 
several  silicates.  When  kept  for  a  considerable  length  of  time  at 
a  temperature  insufficient  to  render  it  perfectly  fluid,  some  of  its 
components  crystallize  out,  the  glass  becomes  opaque,  and  "de- 
vitrification" is  said  to  have  occurred.  The  word  "crystal" 
popularly  applied  to  glass  is  thus  definitely  misleading. 

So-called  quartz-glass  is  made  of  fused  silica  (p.  428). 


STRONTIUM 


495 


Calcium-ion  Ca++:  Analytical  Reactions.  —  Ionic  calcium 
is  colorless.  It  is  bivalent,  and  combines  with  negative  ions. 
Many  of  the  resulting  salts  are  more  or  less  insoluble  in  water. 
Upon  the  insolubility  of  the  carbonate,  phosphate,  and  oxalate  are 
based  tests  for  calcium-ion  in  qualitative  analysis  (see  p.  538). 
The  presence  of  the  element  is  most  easily  recognized  by  the 
brick-red  color  its  compounds  confer  on  the  Bunsen  flame,  and 
by  two  bands  — a  red  and  a  green  one  — which  are  shown  by 
the  spectroscope. 

STRONTIUM  Sr 

The  compounds  of  strontium  resemble  closely  those  of  calcium, 
both  in  physical  properties  and  in  chemical  behavior. 

Occurrence.  — The  carbonate  of  strontium  SrCO3  is  found  as 
strontianite  (Strontian,  a  village  in  Argyleshire).  The  sulphate, 
celestite  SrSO4,  is  more  plentiful.  The  metal  may  be  isolated  by 
electrolysis  of  the  molten  chloride. 

Compounds  of  Strontium.  —  The  compounds  are  all  made 
from  the  natural  carbonate  or  sulphate.  The  former  may  be  dis- 
solved directly  in  acids,  and  the  latter  is  first  reduced  by  means 
of  carbon  to  the  sulphide,  and  then  treated  with  acids. 

Strontium  chloride  SrCl2,6H20,  made  in  one  of  the  above  ways, 
is  deposited  from  solution  as  the  hexahydrate.  The  nitrate 
Sr(N03)2  comes  out  of  hot  solutions  in  octahedrons  which  are 
anhydrous.  From  cold  water  the  tetrahydrate  Sr(N03)2,4H2O  is 
obtained.  The  anhydrous  nitrate  is  mixed  with  sulphur,  charcoal, 
and  potassium  chlorate  to  make  -red  fire."  The  oxide  SrO  may 
be  secured  by  igniting  the  carbonate,  but  it  is  obtained  with 
greater  difficulty  than  is  calcium  oxide  from  calcium  carbonate. 
It  is  generally  made  by  heating  the  nitrate  or  hydroxide. 

Strontium  hydroxide  Sr(OH)2  is  made  by  heating  the  carbonate 
in  a  current  of  superheated  steam: 

SrC03  +  H20  ->  Sr(OH)2  +  C02. 

This  action  takes  place  more  easily  than  does  the  mere  disso- 
ciation of  the  carbonate,  because  the  formation  of  the  hydroxide 
liberates  energy,  and  this  partially  compensates  for  the  energy 


496  COLLEGE    CHEMISTRY 

which  has  to  be  provided  to  decompose  the  carbonate.  The 
lowering  of  the  partial  pressure  of  the  carbon  dioxide  by  the 
steam  also  contributes  to  the  result  (cf.  pp.  476-477).  A  hydrate 
Sr(OH)2,8H20  crystallizes  from  water. 

Strontium-ion  Sr++  is  bivalent,  and  gives  insoluble  compounds 
with  carbonate-ion,  sulphate-ion,  and  oxalate-ion.  The  presence 
of  strontium  is  recognized  by  the  carmine-red  color  which  its 
compounds  give  to  the  Bunsen  flame  (see  also  p.  498).  Its 
spectrum  shows  several  red  bands  and  a  very  characteristic  blue 
line. 

BARIUM  Ba 

The  physical  and  chemical  properties  of  the  compounds  of 
barium  recall  those  of  strontium  and  calcium.  All  the  com- 
pounds of  barium  which  are  soluble  in  water,  or  can  be  brought 
into  solution  by  the  weak  acids  of  the  .digestive  fluids,  are  poison- 
ous. 

Occurrence.  —  Like  strontium,  barium  is  found  in  the  form  of 
the  carbonate,  witherite  BaCOs,  and  the  sulphate  BaSC>4,  heavy 
spar  or  barite  (Gk.  fiapvs,  heavy).  The  free  metal,  which  is  silver- 
white,  may  be  obtained  by  electrolysis  of  the  molten  chloride. 

The  compounds  are  made  by  treating  the  natural  carbonate 
with  acids  directly,  or  by  first  reducing  the  sulphate  with  carbon 
to  sulphide,  or  converting  the  carbonate  into  oxide,  and  then 
treating  the  products  with  acids. 

Compounds  of  Barium.  —  The  precipitated  form  of  barium 
carbonate  BaCOs  is  made  by  adding  sodium  carbonate  to  the 
aqueous  extract  from  crude  barium  sulphide  (q.v.).  Barium  car- 
bonate demands  so  high  a  temperature  (about  1500°)  for  the  at- 
tainment of  a  sufficient  dissociation  tension,  that  special  means 
is  employed  for  its  decomposition.  It  is  heated  with  powdered 
charcoal  (cf.  p.  385) : 

BaC03  +  C  -»  BaO  +  2CO. 

Natural  barium  sulphate  BaSO4  is  the  source  of  many  of  the 
compounds.  The  precipitated  sulphate,  made  by  adding  sul- 
phuric acid  to  the  aqueous  extract  from  barium  sulphide,  is  used  in 
making  white  paint  ("permanent  white"),  in  filling  paper  for 


BARIUM 


glazed  cards,  and  sometimes  as  an  adulterant  of  white  lead    A 
mixture  of  barium  sulphate  and  zinc  sulphide  ZnS,  prepared  i 
special  way,  is  called  lithopone: 

BaS  +  ZnS04  -»  BaS04  j  +  ZnS| . 

Made  into  paint  it  has  greater  covering  power  than  white  lead, 
does  not  darken  with  hydrogen  sulphide  as  does  the  latter,  and  is 
non-poisonous.  Barium  sulphate  is  highly  insoluble  in  water  and 
is  hardly  at  all  affected  by  aqueous  solutions  of  any  chemic 

^Barium  sulphide  BaS,  like  the  sulphides  of  calcium  and  stron- 
tium (p.  273),  is  very  slightly  soluble  in  water,  but  slowly  passes 
into  solution  owing  to  hydrolysis  and  formation  of  the  hydroxide 
and  hydrosulphide.  It  is  made  by  heating  the  pulverized  sul- 
phate with  charcoal :  BaS04  +  4C  ->  BaS  +  4CO. 

Barium  chloride  BaCl2,2H2O  is  manufactured  by  heating  the 
sulphide  with  calcium  chloride.  The  whole  treatment  of  the 
heavy  spar  is  carried  out  in  one  operation: 

BaS04  +  4C  +  CaCl2  ->  4CO  +  BaCl2  +  CaS. 

By  rapid  extraction  with  water,  the  chloride  can  be  separated 
from  the  calcium  sulphide  before  much  decomposition  of  the 
latter  (cf.  p.  461)  has  taken  place. 

Barium  chlorate  Ba(ClO3)2  is  made  by  treating  the  precipitated 
barium  carbonate  with  a  solution  of  chloric  acid.  It  is  deposited 
in  beautiful  monoclinic  crystals,  and  is  used  with  sulphur  and 
charcoal  in  the  preparation  of  "green  fire." 

Barium  monoxide  BaO  is  manufactured  from  the  carbonate  (see 
above)  or,  in  pure  form,  by  heating  the  nitrate.  The  oxide  unites 
vigorously  with  water  to  form  the  hydroxide. 

The  monoxide,  when  heated  in  a  stream  of  air  or  oxygen,  gives 
barium  peroxide:  2BaO  +  02  *±  2BaO2,  as  a  compact  gray  mass. 
This  change  and  its  reversal  constitute  the  basis  of  Brm  s  process 
for  obtaining  oxygen  from  the  air.  At  a  suitable,  high  temper- 
ature  the  air  is  forced  in  under  pressure,  causing  the  action  to  go 
forward,  while  the  nitrogen  escapes  by  a  valve  at  the  far  end  ot 
the  apparatus.  Then,  without  change  of  temperature,  by  re- 
versing the  pumps,  oxygen  is  taken  out,  and  the  reaction  goes 
backwards.  This  alternation  makes  the  process  a  continuous 


498  COLLEGE    CHEMISTRY 

one.  A  hydrate,  BaO2,8H2O,  is  thrown  down  as  a  crystalline 
precipitate  when  hydrogen  peroxide  solution  is  added  to  a  solu- 
tion of  barium  hydroxide : 

Ba(OH)2  +  H202  <=*  BaO2 1  +  2H2O. 

Barium  peroxide  is  used  in  the  manufacture  of  hydrogen  peroxide. 

Barium  hydroxide  Ba(OH)2,  is  made  by  union  of  the  oxide  with 
water,  or  by  leading  moist  carbon  dioxide  over  the  sulphide  and 
decomposing  the  resulting  carbonate  with  superheated  steam 
(p.  495).  It  is  the  most  soluble  of  the  hydroxides  of  this  group, 
and  gives,  therefore,  the  highest  concentration  of  hydroxide- 
ion.  The  solution  is  known  as  "  baryta- water."  It  is  also  the 
most  stable  of  the  three  hydroxides,  and  may  be  melted  with- 
out decomposition.  A  hydrate  Ba(OH)2,8H2O  crystallizes  from 
water. 

Barium  nitrate  Ba(NOs)2  is  made  by  the  action  of  nitric  acid  on 
the  sulphide,  oxide,  hydroxide,  or  carbonate  of  barium.  The 
crystals  from  aqueous  solution  are  anhydrous. 

Analytical  Reactions  of  the  Calcium  Family.  —  Barium-ion 
Ba++  is  a  colorless,  bivalent  ion.  Many  of  its  compounds  are  in- 
soluble in  water,  and  the  sulphate  is  insoluble  in  acids  also.  The 
spectrum  given  by  the  salts  contains  a  number  of  green  and  orange 
lines. 

In  solutions  of  salts  of  calcium,  strontium,  and  barium,  the  ions 
may  be  distinguished  by  the  fact  that  calcium  sulphate  solution 
will  precipitate  the  strontium  and  barium  as  sulphates,  but  will 
leave  salts  of  calcium  in  dilute  solution  unaffected.  Similarly, 
strontium  sulphate  solution  precipitates  barium  sulphate,  and 
does  not  give  any  result  with  salts  of  the  first  two.  The  oxalate 
of  calcium  is  precipitated  in  presence  of  acetic  acid,  while  the 
oxalates  of  strontium  and  barium  are  not  (cf.  p.  484),  and  there 
are  other  differences  of  a  like  nature  in  the  solubilities  of  the  salts. 

Exercises.  —  1.  Arrange  the  chromates  of  the  metals  of  this 
family  in  the  order  of  solubility  (see  Table).  Compare  the  solubili- 
ties with  those  of  the  carbonates,  oxalates,  and  sulphates  of  the 
metals  of  the  same  family. 

2.   What  is  meant  by  fluorescence  (cf.  any  book  on  physics)? 


BARIUM  499 

3.  What  will  be  the  ratio  by  volume,  at  150°,  of  the  nitrogen 
peroxide  and  oxygen  given  off  by  the  decomposition  of  calcium 
nitrate?    What  would  be  the  nature  of  the  difference  between  the 
ratio  at  150°  and  that  at  room  temperature  (cf.  p.  352)? 

4.  Apply  the  rule  of  precipitation  to  the  case  of  adding  sodium 
carbonate  to  a  solution  of  barium  chloride. 

5.  Explain  in  terms  of  the  ionic  hypothesis  the  precipitation  of 
the  sulphate  of  strontium  by  calcium  sulphate  solution,  and  the 
absence  of  precipitation  when  the  latter  is  added  to  a  dilute  solu- 
tion of  a  soluble  salt  of  calcium. 

6.  What  inference  do  you  draw  from  the  fact  that  the  oxalates 
of  barium  and  strontium  are  not  precipitated  in  presence  of  acetic 
acid,  while  the  oxalate  of  calcium  is  so  precipitated?     Is  the  infer- 
ence  confirmed  by  reference  to  the  solubility  data? 

7.  Explain  the  fact  that  strontium  and  calcium  chromates  are 
easily  dissolved  by  acetic  acid,  while  barium  chromate  is  dissolved 
only  by  active  mineral  acids. 

8.  Explain  the  fact  that  all  the  carbonates,   save  those  of 
sodium,  potassium,  and  thallium,  are  precipitated  in  neutral  solu- 
tions, but  not  in  acidified  solutions.     Why  is  the  precipitation 
incomplete  when  carbon  dioxide  is  led  through  solutions  of  salts 
of  the  metals,  but  more  complete  when  the  hydroxides  of  the 
metals  are  used? 

9.  Construct  a  table  for  the  purpose  of  comparing  the  proper- 
ties of  the  free  elements  of  this  family  and  also  the  properties  of 
their  corresponding  compounds. 

10.  Are  the  elements  of  this  family  typical  metals  (p.  436)? 


CHAPTER  XXXVII 
COPPER,    SILVER,    GOLD 

THE  three  metals  of  this  family,  being  found  free  in  nature,  are 
amongst  those  which  were  known  in  early  times.  They  are  the 
metals  universally  used  for  coinage  and  for  ornamental  purposes. 
They  are  the  three  best  conductors  of  electricity  (p.  436). 

The  Chemical  Relations  of  the  Copper  Family.  —  Copper 
(Cu,  at.  wt.  63.6),  silver  (Ag,  at.  wt.  107.88),  and  gold  (Au,  at.  wt. 
197.2)  occupy  the  right  side  in  the  second  column  of  the  table  of 
the  periodic  system,  and  the  chemical  relations  of  these  elements 
are  in  many  ways  in  sharp  contrast  to  those  of  the  alkali  metals, 
their  neighbors,  on  the  left  side: 

ALKALI  METALS  COPPER,  SILVER,  GOLD 

Very  active;  rapidly  oxidized  by  air;      Amongst  least  active  metals;    only 

displace  all  other  metals  from  com-          copper  is  oxidized  by  air;  displaced 

bination  (E.-M.  series,  p.  60).  by  most  other  metals. 

All  univalent  and  give  but  one  series      Cu1  and  Cu11:  two  series.    Ag1:  one 

of  compounds.     Halides  all  soluble          series.     Au1  and  Aum:  two  series. 

in  water.  Chlorides  of  univalent  series  insol. 

Oxides  and  hydroxides  strongly  basic,      Oxides  and  hydroxides  feebly  basic 

and  halides  not  hydrolyzed  (p.  437).          (except  Ag2O) ;  halides  hydrolyzed 

(except  Ag-halides).     Hence,  basic 
salts  are  numerous. 

Never  found  in  anion.    Give  no  com-      Frequently  in  anion,  e.g.,  K.Cu(CN)2, 
plex  cations.  KAg(CN)2,  KAuO2,  K.Au(CN)2, 

and  also  in  complex  cations,  e.g., 
Ag(NH3)2.OH  and  Cu(NH3)4.(OH)2 

On  account  of  their  inactivity  towards  oxygen,  and  their  easy 
recovery  from  combination  by  means  of  heat,  silver  and  gold, 
together  with  the  platinum  family,  are  known  as  the  "noble 
metals." 

500 


COPPER  501 

COPPER  Cu 

Chemical  Relations  of  the  Element.  —  Copper  is  the  first 
metallic  element  showing  two  valences  which  we  have  encountered. 
In  such  cases  two  more  or  less  complete,  independent  series  of  salts 
are  known.  These  are  here  distinguished  as  cuprous  (univalent) 
and  cupric  (bivalent)  salts.  The  methods  by  which  a  compound  of 
one  series  may  be  converted  into  the  corresponding  compound  of 
the  other  series  should  be  noted. 

The  chief  cuprous  compounds  are  Cu2O,  CuCl,  CuBr,  Cul, 
CuCN,  Cu2S.  The  cuprous  compound  is  in  each  of  these  cases 
more  stable  (p.  93)  than  the  corresponding  cupric  compound,  and 
is  formed  from  the  latter  either  by  spontaneous  decomposition, 
as  in  the  cases  of  the  iodide  and  cyanide  (2CuI2  — »  2CuI  +  I2),  or 
on  heating.  The  cuprous  halides  and  cyanide  are  colorless  and 
insoluble  in  water.  Cuprous-ion  Cu+  seems  to  be  colorless.  The 
cuprous  salts  of  oxygen  acids  have  few  applications. 

The  cupric  compounds  are  more  numerous,  as  they  include  also 
stable  and  familiar  salts  of  oxygen  acids,  like  CuSO4,  Cu(NOs)2,  etc. 
The  anhydrous  salts  are  usually  colorless  or  yellow,  but  cupric-ion 
Cu++  is  blue,  and  so,  therefore,  are  the  aqueous  solutions  of  the 
salts.  The  cupric  are  more  familiar  than  the  cuprous  compounds, 
since  cupric  oxide,  sulphate,  and  acetate  are  the  compounds  of 
copper  which  most  frequently  find  employment  in  chemistry  and 
in  the  arts.  All  the  soluble  salts  of  copper  are  poisonous. 

In  addition  to  (1)  having  two  valences  Cu1  and  Cu11,  and  there- 
fore two  series  of  compounds  (two  oxides,  two  chlorides,  etc.),  each 
of  these  states  of  copper  also  joins  with  other  elements  to  form  (2) 
complex  positive  ions  such  as  Cu(NH3)2+  and  Cu(NH3)4++,  just  as 
hydrogen  and  ammonia  form  the  complex  positive  ion  NH4+,  and 
the  univalent  form  also  gives  (3)  stable  complex  negative  ions  such 
as  Cu(CN)2~,  CuCl2~.  None  of  the  metallic  elements  discussed  in 
the  two  preceding  chapters  showed  any  of  these  peculiarities. 
Many  of  the  metals  to  be  discussed  later  exhibit  one  or  more  of 
them,  however.  Especial  attention  should  therefore  be  given  to  the 
chemistry  of  copper,  in  order  that  the  behavior  which  such  relations 
entail  may  be  mastered  at  the  first  encounter,  and  the  same  rela- 
tions may  be  instantly  recognized  and  understood  when  they 
reappear  in  other  connections. 


502  COLLEGE    CHEMISTRY 

There  is  only  one  other  peculiarity  which  a  metallic  element  fre- 
quently shows,  although  copper  does  not  exhibit  it.  This  is  (4)  the 
ability  of  its  hydroxide  to  be,  not  only  basic,  as  metallic  hydroxides 
by  definition  (p.  436)  must  be,  but  also  acidic.  This  behavior  we 
encounter  first  in  the  case  of  gold  (see  p.  520)  and  in  simpler  and 
more  familiar  form  in  the  case  of  zinc  (see  next  chapter). 

Occurrence.  —  Copper  is  found  free  in  the  Lake  Superior 
region.  The  sulphides,  copper  pyrites  CuFeS2  and  chalcocite 
Cu2S,  are  worked  in  Montana,  Utah,  southwest  England,  Spain, 
and  Germany.  Malachite,  Cu2(OH)2CO3  (=  Cu(OH)2,CuCO3),  a 
basic  carbonate,  is  mined  in  Arizona,  Siberia,  and  elsewhere. 
Cuprite  or  ruby  copper  Cu20  is  also  an  important  ore. 

Extraction  from  Ores.  —  For  isolating  native  copper  it  is 
only  necessary  to  separate  the  metal,  by  grinding  and  washing, 
from  the  rock  through  which  it  ramifies,  and  to  melt  the  almost 
pure  powder  of  copper  with  a  flux  (p.  438).  The  carbonate  and 
oxide  ores  require  coal,  in  addition,  for  the  removal  of  the  oxygen. 

The  liberation  of  copper  from  the  sulphide  ores  is  difficult,  and 
often  involves  very  elaborate  schemes  of  treatment.  This  arises 
from  the  fact  that  many  copper  ores  contain  a  large  amount  of  the 
sulphides  of  iron,  and  these  have  to  be  removed  by  conversion  into 
oxide  (by  roasting)  and  then  into  silicate  (with  sand).  The  silicate 
forms  a  flux,  and  separates  itself  from  the  molten  mixture  of  copper 
and  copper  sulphide.  In  Montana  it  is  found  possible  to  abbrevi- 
ate the  treatment.  The  ore  is  first  roasted  until  partially  oxidized. 
It  is  then  melted  in  a  cupola  or  a  reverberatory  furnace,  and  placed 
in  large  iron  vessels  like  Bessemer  converters  (q.v.)  provided  with  a 
lining  rich  in  silica.  A  blast  of  air  mixed  with  sand  is  next  blown 
through  the  mass.  The  iron  is  completely  oxidized  to  FeO  and 
made  into  silicate  FeSi03,  the  sulphur  escapes  as  sulphur  dioxide, 
and  arsenic  and  lead  are  likewise  removed  by  this  treatment.  The 
silicate  of  iron  floats  as  a  slag  upon  the  copper  when  the  contents  of 
the  converter  are  poured  out.  The  resulting  copper  is  pure  enough 
to  be  cast  in  large  plates  and  purified  by  electrolysis  (see  p.  511). 

Much  copper  ore  is  of  low  grade,  containing  perhaps  only  2  per 
cent  of  copper  ore  and  98  per  cent  of  rock  material.  From  such 
ores  the  usual  methods  of  washing  often  recover  only  70  per  cent 


COPPER  503 

or  less  of  the  copper  ore  present,  and  30  per  cent  or  more  is  lost. 
The  froth  flotation  process  raises  the  proportion  recovered  to  85  or 
90  per  cent  of  the  whole.  The  finely  crushed  ore  is  agitated  with 
water,  to  which  is  added  some  cheap  oil  and  sometimes  a  little 
sulphuric  acid.  The  mixture  is  then  allowed  to  flow  into  a  larger 
tank  of  water,  in  which  the  rock  material  immediately  sinks  to  the 
bottom  while  the  particles  of  ore  are  contained  in  the  oily  froth 
which  rises  to  the  top.  The  plant  also  occupies  less  than  one-tenth 
of  the  space,  and  uses  less  than  half  the  power  required  for  treating 
the  same  amount  of  ore  by  washing. 

The  world's  production  (1913)  is  about  a  million  metric  tons,  of 
which  the  United  States  furnished  58  per  cent,  South  America  11, 
Japan  6,  and  Germany  4. 

Physical  and  Chemical  Properties. — Copper  is  red  by  re- 
flected and  greenish  by  transmitted  light.  It  melts  at  1083°,  and 
therefore  much  more  easily  than  pure  iron  (1530°).  Sp.  gr.  8.93. 

In  ordinary  air  copper  becomes  slowly  covered  with  a  green  basic 
carbonate  (not  verdigris,  q.v.).  It  does  not  decompose  water  at 
any  temperature  or  displace  hydrogen  from  dilute  acids  (p.  60). 
The  metal  attacks  oxygen  acids  (pp.  275,  354),  however.  Sea- 
water  and  air  slowly  corrode  the  copper  sheathings  of  ships,  giving 
the  basic  chloride  Cu4(OH)6Cl2,H2O(=  3Cu(OH)2,CuCl2,H20), 
which  is  found  in  nature  as  atakamite. 

On  account  of  its  resistance  to  the  action  of  acids,  copper  is  used 
for  many  kinds  of  vessels,  for  covering  roofs  and  ships7  bottoms, 
and  for  coins.  It  furnishes  also  electrotype  reproductions  of 
medals,  of  engraved  plates,  of  type,  etc.  (see  p.  510).  Great  quan- 
tities of  the  metal  are  used  in  electrical  plants  and  appliances. 

Alloys.  —  The  qualities  of  copper  are  modified  for  special  pur- 
poses by  alloying  it  with  other  metals.  Brass  contains  18-40  per 
cent  of  zinc,  and  melts  at  a  lower  temperature  (p.  134)  than  does 
copper.  A  variety  with  little  zinc  is  beaten  into  thin  sheets,  giving 
Dutch  metal  ("gold  leaf").  Bronze  contains  3-8  per  cent  of  tin, 
11  or  more  per  cent  of  zinc,  and  some  lead.  It  was  used  for  making 
weapons  and  tools  before  means  of  hardening  iron  were  known,  and 
later,  on  account  of  its  fusibility,  continued  to  be  employed  for 
castings  until  displaced  largely  by  cast  iron  (discovered  in  the  eight- 


504  COLLEGE    CHEMISTRY 

eenth  century).  Gun  metal  contains  10  per  cent,  and  bell  metal 
20-24  per  cent  of  tin.  German  silver  contains  19-44  per  cent  of 
zinc  and  6-22  per  cent  of  nickel,  and  shows  none  of  the  color  of 
copper.  In  many  of  these  alloys  the  metals  are  partly  in  the  form 
of  chemical  compounds,  such  as  Cu3Sn  and  Cu2Zn3. 

Cupric  Chloride  CuCl^H^O. —  This  compound  is  made  by 
union  of  copper  and  chlorine,  or  by  treating  the  hydrate  or  car- 
bonate'with  hydrochloric  acid.  The  blue  crystals  of  a  hydrate, 
CuCl2,2H2O,  are  deposited  by  the  solution.  The  anhydrous  salt 
is  yellow.  Dilute  solutions  are  blue,  the  color  of  cupric-ion,  but 
concentrated  solutions  are  green  on  account  of  the  presence  of  the 
yellow  molecules  (p.  249) .  The  aqueous  solution  is  acid  in  reaction 
(p.  437).  When  excess  of  ammonium  hydroxide  is  added  to  the 
solution,  the  basic  chloride,  cupric  oxychloride  Cu4(OH)6Cl2  (see 
above),  which  is  at  first  precipitated,  redissolves,  and  a  deep-blue 
solution  is  obtained  (see  p.  506).  This  on  evaporation  yields  deep- 
blue  crystals  of  hydrated  ammonio-cupric  chloride  Cu(NH3)4.Cl2, 
H20.  The  deep-blue  color  of  the  solution,  which  is  given  by  all 
cupric  salts,  is  that  of  ammonio-cupric-ion  Cu(NH3)4++.  The  dry 
salt  also  absorbs  ammonia,  giving  CuCl2,6NH3,  but  a  reduction  of 
pressure  results  finally  in  the  loss  of  all  the  ammonia. 

Cuprous  Chloride  CuCl.  —  It  may  be  made  by  boiling  cupric 
chloride  solution  with  hydrochloric  acid  and  copper  turnings: 

CuCl2  +  Cu->2CuCl,    or    Cu++  +  Cu  -*  2Cu+. 

The  solution  contains  compounds  of  cuprous  chloride  with  hydrogen 
chloride  HCl,CuCl,  or  HCuCl2  and  H2CuCl3,  which  are  decomposed 
when  the  acid  solution  is  diluted  with  water.  The  cuprous  chlo- 
ride is  insoluble  in  water,  and  forms  a  white  crystalline  precipitate. 

The  foregoing  action  is  an  illustration  of  the  fifth  kind  of  ionic 
chemical  change,  namely,  that  in  which  a  change  in  valence  (and  also 
in  the  amount  of  the  electrical  charge),  occurs,  without  any  altera- 
tion in  the  composition  of  the  ionic  substance.  For  other  illustra- 
tions see  pp.  158  (Mn++++  +  4Q-  _>  Mn++  +  2CP  +  C12),  501. 

Cuprous  chloride  is  hydrolyzed  quickly  by  hot  water,  giving, 
finally,  red,  hydrated  cuprous  oxide,  Cu20.  When  dry  it  is  not 
affected  by  light,  but  in  the  moist  state  becomes  violet  and,  finally, 


COPPER  505 

nearly  black.  The  action  is  said  to  be  2CuCl  — >  CuCl2  +  Cu.  Jn 
moist  air  it  turns  green,  and  is  oxidized  to  cupric  oxychloride  (p. 
504).  It  is  dissolved  by  hydrochloric  acid,  giving  the  colorless 
complex  acids  HCuCl2  and  H2CuCl3  just  mentioned  (see  below). 
The  solution  is  oxidized  by  the  air,  turning  first  brown  and  then 
green,  and  finally  depositing  the  cupric  oxychloride.  It  also  has 
the  power  of  absorbing  carbon  monoxide,  to  form  a  compound 
said  to  be  Cu(CO)Cl,H2O,  and  the  property  is  used  to  separate 
this  gas  in  analyzing  mixtures  of  gases.  Cuprous  chloride  also 
dissolves  (see  p.  506)  in  ammonium  hydroxide,  giving  ammonio 
cuprous  chloride  Cu(NH3)2.Cl,  the  ion  Cu(NH3)2+  being  colorless. 
The  solution  is  quickly  oxidized  by  the  air,  turns  deep-blue,  and 
then  contains  Cu(NH3)4++. 

The  Bromides  and  Iodides  of  Copper.  —  By  treatment  of 
copper  with  bromine-water,  and  slow  evaporation  of  the  solution, 
jet-black  crystals  of  anhydrous  cupric  bromide  CuBr2  are  obtained 
(cf.  p.  249).  When  dry  cupric  bromide  is  heated,  bromine  is  given 
off,  and  cuprous  bromide  CuBr  remains. 

Cupric  iodide  CuI2  appears  to  be  unstable  at  ordinary  tempera- 
tures. When  a  soluble  iodide  is  added  to  a  cupric  salt,  a  white 
precipitate  of  cuprous  iodide  Cul  and  free  iodine  are  obtained: 

2Cu++  +  41"  *=¥  2CuI  J,  +  I2. 

The  Solution  of  Insoluble  Salts  when  Complex  Ions  are 
Formed.  —  The  solution  of  an  insoluble  salt  like  cuprous  chloride 
by  hydrochloric  acid  or  ammonium  hydroxide  is  typical  of  a  great 
variety  of  actions  of  which  we  here  meet  with  the  first  examples. 
Compound  or  complex  ions  are  formed  (cf.  p.  501).  The  explana- 
tion involves  only  principles  already  used  in  other  cases. 

The  dissolving  of  cuprous  chloride  in  hydrochloric  acid  (p.  505), 
to  form  soluble,  complex,  highly  ionized  acids  like  H.CuCl2  is  a 
typical  case.  The  complex  negative  ion  CuCl2~  which  is  formed 
is  very  little  dissociated  (CuCl2~<=^  Cu++ 2C1~),  and  gives  a 
smaller  concentration  of  Cu+  than  does  the  insoluble  cuprous  chlo- 
ride. The  ion-product  of  cuprous  chloride,  and  the  concentra- 
tion relations  of  the  ionic  substance  CuCl2~  and  its  dissociation 
products  (Cu+  and  2C1~)  are  symbolized  as  follows: 

rrn+i  v  rrn     K'       [Cu+1  x  [cr]2     K 

[Cu+]  X  [Cl  ]  =  K  K. 


506  COLLEGE    CHEMISTRY 

The  value  of  [Cu+]  from  cuprous  chloride  (first  formula)  is,  in 
general,  greater  than  its  value  from  the  ion  CuCl2~  of  HCuCl2 
(second  formula),  when  excess  of  HC1  is  present.  Hence,  the  Cu+ 
tends  to  pass  over  into  the  more  stable  compound,  where  it  is  more 
completely  combined.  More  CuCl  dissolves  to  replace  the  Cu+ 
which  has  been  removed,  and  the  change  stops  when  the  CuCl  is 
all  dissolved,  or  the  values  of  [Cu+]  from  both  compounds  have 
become  equal.  Thus,  the  complex  ion  is  formed  at  the  expense  of 
the  Cu+  of  the  insoluble  cuprous  chloride,  and  the  latter  goes  into 
solution  progressively  in  the  effort  to  restore  the  balance: 

CuCl  (solid)  fc?  CuCl  (ddvd)±*Cl-  +Cu+  UCuC1-/dslvcn 
2HC1  ^2H+  +  2Crj<- 

The  same  exact  laws  of  equilibrium  used  in  discussing  the  dissolv- 
ing of  salts  by  acids  (p.  481)  may  be  applied  to  the  whole  procedure. 

The  dissolving  of  cuprous  chloride  by  the  free  ammonia  of  ammo- 
nium hydroxide  is  explained  in  the  same  way.  The  only  difference 
is  that  here  the  copper  is  in  a  complex  positive  ion.  The  ion 
Cu(NH3)2"f  gives  little  Cu+  —  less  than  does  cuprous  chloride,  in 
spite  of  the  insolubility  of  the  latter.  Hence  the  salt  passes  into 
solution  until  the  ion-product  [Cu+]  X  [Cl~],  with  continually  in- 
creasing [Cl~],  reaches  the  value  for  a  saturated  solution  or  until 
the  solid  is  exhausted. 

The  deep-colored  ion  Cu(NH3)4++  given  by  cupric  chloride  and 
other  cupric  salts  is  also  very  little  ionized.  Hence  ammonium 
hydroxide  dissolves  all  the  insoluble  cupric  compounds  save  only 
cupric  sulphide,  which  is  the  most  insoluble  of  all  —  that  is,  the  one 
giving  the  smallest  concentration  of  cupric-ion.  Conversely,  the 
sulphide  is  the  only  insoluble  compound  of  copper  which  can  be 
precipitated  from  ammoniacal  solution. 

Foregoing  Explanation  Restated.  —  We  may  restate  the 
explanation  by  answering  a  question:  Why  does  cuprous  chloride 
interact  with,  and  go  into  solution  in  hydrochloric  acid?  Because 
it  forms  a  complex  compound  HCuCl2,  and,  with  the  concentrations 
usually  employed,  the  molecular  concentration  of  cuprous-ion  in  the 
solubility  product  of  cuprous  chloride  is  greater  than  the  molecular 
concentration  of  the  same  ion  in  the  solution  of  the  complex  com- 
pound. 


COPPER  507 

The  answer  in  other  cases  takes  the  same  form.  Thus,  for 
cupric  hydroxide  Cu(OH)2  dissolving  in  ammonium  hydroxide 
solution,  substitute  cupric  hydroxide  for  cuprous  chloride  and 
Cu(NH3)4(OH)2  for  HCuCl2. 

Cuprous  Oxide  Cu%O.  —  This  oxide  is  red  in  color,  and  natural 
specimens  show  octahedral  forms.  It  is  produced  by  oxidation  of 
finely  divided  copper  at  a  gentle  heat,  or  by  the  addition  of  bases  to 
cuprous  chloride,  and  is  best  made  by  the  action  of  glucose  (p.  404) 
on  cupric  hydroxide  (see  Fehling's  solution,  below).  The  simple 
hydroxide,  CuOH,  is  unknown,  but  the  above  mentioned  pre- 
cipitate is  a  hydrated  oxide  4Cu2O,H20,  and  yields  Cu20  when 
heated. 

Cuprous  oxide  is  acted  upon  by  hydrochloric  acid,  giving  cu- 
prous chloride,  or  rather  HCuCl2.  It  also  dissolves  in  ammonium 
hydroxide,  giving,  probably,  Cu(NH3)2.OH,  which  is  colorless. 

Cupric  Oxide  and  Hydroxide.  —  Cupric  oxide  CuO  (black)  is 
formed  by  heating  copper  in  a  stream  of  oxygen  or,  in  less  pure 
form,  by  igniting  the  nitrate,  carbonate,  or  hydroxide.  When 
heated  strongly  it  loses  some  oxygen,  and  is  partly  reduced  to 
Cu2O. 

Cupric  hydroxide  Cu(OH)2  is  precipitated  as  a  gelatinous  sub- 
stance by  addition  of  sodium  or  potassium  hydroxide  to  a  solution 
of  a  cupric  salt:  Cu++  +  20H~  -*  Cu(OH)2.  When  the  mixture 
is  boiled,  the  hydroxide  loses  water  and  forms  a  black  hydrated 
cupric  oxide  Cu(OH)2,2CuO(?). 

The  hydroxide  interacts  with  ammonium  hydroxide,  forming  the 
soluble  compound  Cu(NH3)4-(OH)2,  which  has  a  deep-blue  color. 
Cellulose  (cotton  or  paper)  is  soluble  in  this  solution,  and  is  re- 
precipitated  by  sulphuric  acid.  Artificial  silk  is  made  by  pressing 
the  solution  through  dies  into  the  precipitant.  Paper  and  cotton 
goods,  when  passed  first  through  one  and  then  the  other  of  these 
liquids,  receives  a  tough,  waterproof  surface. 

Cupric  hydroxide  interacts  with  a  solution  of  Rochelle  salt, 
giving  a  soluble  compound.  The  liquid  is  known  as  Fehling's 
solution,  and  is  used  in  testing  for,  and  estimating  quantities  of 
glucose  (p.  404),  and  other  reducing  substances.  Cuprous  oxide 
is  precipitated  (see  above). 


508  COLLEGE    CHEMISTRY 

Cupric  Nitrate  Cu(]YO3)2,6/f2O.  —  The  nitrate  is  made  by 
treating  cupric  'oxide  or  copper  with  nitric  acid  (p.  354),  and  is 
obtained  from  the  solution  as  a  deliquescent,  crystalline  hexahy- 
drate.  When  dehydrated  at  65°  the  salt  is  partly  hydrolyzed, 
and  a  basic  nitrate  Cu4(OH)6(N03)2  remains. 

Carbonate  of  Copper.  —  No  normal  carbonate  (CuC03)  can  be 
obtained.  A  basic  carbonate  (malachite)  is  found  in  nature,  and  is 
precipitated  by  adding  soluble  carbonates  to  cupric  salts: 

2CuS04  +  2Na2C03  +  H20  -*  Cu2(OH)2CO3  +  2Na2S04  +  CO2. 


The  carbonate,  if  formed,  would  be  hydrolyzed  by  water  (p.  437). 

Cyanides  of  Copper.  —  With  potassium  cyanide  and  a  solution 
of  a  cupric  salt,  cupric  cyanide  Cu(NC)2  is  precipitated.  This  is 
not  stable,  however,  and  gives  off  cyanogen,  leaving  cuprous 
cyanide  CuNC: 

2Cu(NC)2  -»  2CuNC  +  C2N2. 

Cuprous  cyanide  is  insoluble  in  water,  but  interacts  with  an  excess 
of  potassium  cyanide  solution,  producing  a  colorless  liquid,  from 
which  K.Cu(CN)2  (=  KCN,CuCN),  potassium  cuprocyanide,  may 
be  obtained  in  colorless  crystals.  The  complex  anion  Cu(CN)2~  is 
so  little  ionized  to  Cu+  and  2CN~  that  all  insoluble  copper  com- 
pounds, including  cupric  sulphide,  are  dissolved  by  potassium 
cyanide;  and  none  of  them  can  be  precipitated  from  the  solution. 
Zinc  is  actually  unable  to  displace  copper  from  such  a  solution. 
The  cause  of  the  solution  of  the  salts  is  the  same  as  when  the  com- 
plex ions  Cu(NH3)2+,  Cu(NH3)4++,  and  CuCl2~are  formed  (p.  505). 

Cupric  Acetate.  —  By  the  oxidation  of  plates  of  copper,  sepa- 
rated by  cloths  saturated  with  acetic  acid  (vinegar)  ,  a  basic  acetate 
of  copper  (verdigris)  is  obtained: 

6Cu  +  8HC2H302  +  3O2-+2Cu3(OH)2(C2H302)4  +  2H20. 

It  is  used  in  manufacturing  green  paint,  is  insoluble  in  water,  and  is 
unaffected  by  light.  It  dissolves  in  acetic  acid,  and  green  crystals 
of  the  normal  acetate  Cu(C2H3O2)2,H2O  are  obtained  from  the 
solution.  The  basic  acetate  is  used  in  preparing  Paris  green. 


COPPER  509 

Cupric  Sulphate  CuSO*.  —  This  salt  is  obtained  by  heating 
copper  in  a  furnace  with  sulphur,  and  admitting  air  to  oxidize  the 
cuprous  sulphide.  The  mixture  of  cupric  sulphate  and  cupric  oxide 
which  is  formed  is  treated  with  sulphuric  acid.  The  salt  is  also 
made  by  allowing  dilute  sulphuric  acid  to  trickle  over  granulated 
copper,  while  air  has  free  access  to  the  material: 

2Cu  +  2H2S04  +  02  ->  2CuS04  +  2H20. 

This  is  an  example  of  the  use  of  two  reagents  which  separately  have 
little  or  no  action  (cf.  pp.  426,  431).  When  concentrated  and  very 
hot,  sulphuric  acid  will  itself  act  as  the  oxidizing  agent  (cf.  p.  276). 

Cupric  sulphate  crystallizes  as  pentahydrate  CuSO4,5H20  in  blue 
asymmetric  crystals  (Fig.  52,  p.  95),  and  in  this  form  is  called  blue- 
stone  or  blue  vitriol.  The  aqueous  solution  has  an  acid  reaction 
(p.  437).  The  anhydrous  salt  is  white,  and  can  be  crystallized  in 
thin  needles  from  solution  in  hot,  concentrated  sulphuric  acid  (cf. 
p.  95).  Cupric  sulphate  is  employed  in  copper-plating  (see  p. 
510),  in  batteries,  and  as  a  mordant  in  dyeing  (q.v.).  A  minute 
proportion  is  added  to  drinking  water,  to  destroy  algce,  which  other- 
wise propagate  in  the  reservoirs  and  give  a  disagreeable  taste  and 
odor  to  the  water.  A  solution,  mixed  with  milk  of  lime  (Cu(OH)2 
is  precipitated),  Bordeaux  mixture,  is  largely  used  as  a  spray  on 
grape  vines  and  other  plants  to  prevent  the  growth" of  fungi. 

When  ammonium  hydroxide  is  added  to  cupric  sulphate  solu- 
tion, a  pale-green  basic  sulphate  Cu4(OH)6S04(?)  is  first  precipi- 
tated. With  excess  of  the  hydroxide,  the  blue  Cu(NH3)4++  ion 
(p.  506)  is  formed,  and  crystals  of  ammonio-cupric  sulphate 
Cu(NH3)4.S04,H20  can  be  obtained  from  the  solution. 

Cupric  sulphate  also  combines  with  potassium  and  ammonium 
sulphates,  giving  double  salts  of  the  form  CuSO4,K2SO4,6H20,  which 
are  deposited  in  large,  monosymmetric  crystals  from  the  mixed 
solutions.  Double  salts  (p.  245)  exist  as  such  in  the  solid  form 
only  and,  in  water,  are  resolved  into  the  components  and  their 
ions. 

The  Sulphides  of  Copper.  —  Cuprous  sulphide  Cu2S  occurs  in 
nature  in  rhombic  crystals  of  a  gray,  metallic  appearance.  It  is 
the  sulphide  formed  by  direct  union  of  the  elements.  Cupric 
sulphide  CuS  is  deposited  as  a  black  precipitate  when  hydrogen 


510 


COLLEGE    CHEMISTRY 


sulphide  is  led  through  a  solution  of  a  cupric  salt.  When  heated, 
it  leaves  cuprous  sulphide  and  sulphur  vapor  is  given  off. 

Analytical  Reactions  of  Compounds  of  Copper.  —  The  ion 

of  ordinary  cupric  salts,  cupric-ion  GU++,  is  blue,  and  that  of 
cuprous  salts,  cuprous-ion  Cu+,  is  colorless.  Cuprous  solutions, 
however,  are  easily  oxidized  by  the  air  and  become  blue.  In  solu- 
tions containing  cupric-ion,  hydrogen  sulphide  precipitates  cupric 
sulphide,  even  in  presence  of  acids  (p.  483).  Bases  throw  down 
the  blue  hydroxide,  and  carbonates  precipitate  a  green  basic  salt 
(p.  508).  Potassium  ferrocyanide  gives  the  brown,  gelatinous 
cupric  ferrocyanide: 

2Cu.S04  +  K4.Fe(CN)6^±Cu2.Fe(CN)6 1  +  2K2S04. 

A  characteristic  test  is  the  formation  of  the  deep-blue  Cu(NH3)4++ 
ion  with  excess  of  ammonium  hydroxide.  This  solution  itself 
responds  to  certain  precipitants  only  (e.g.,  H2S).  Solutions  of 
complex  cuprous  cyanides  such  as  K.Cu(CN)2  are  colorless,  and  do 
not  respond  to  any  of  the  above  tests.  With  microcosmic  salt  or 
borax  (pp.  371,  433),  copper  compounds  form  a  bead  which  is 
blue  in  the  oxidizing  part  of  the  flame  and  becomes  red  and  opaque 
(liberation  of  copper)  in  the  reducing  flame. 

Electrotyping.  —  When  plates  of  platinum,  connected  with  a 
battery,   are  immersed  in   cupric   sulphate   solution,    copper   is 

deposited  on  the  cathode  (negative 
pole).  The  sulphate-ion  S04=  mi- 
grates (p.  229)  towards  the  anode 
(positive  pole)  and  there  liberates 
sulphuric  acid  and  oxygen  (p.  228). 
If,  however,  the  anode  is  made  of 
copper,  the  SO4=  migrates,  but  is  not 
discharged.  Instead,  copper  goes  into 
solution  (Fig.  120)  as  Cu++,  in  amount 
equal  to  that  deposited  on  the  other 
pole.  Thus,  the  only  changes  are,  (1)  an  increase  in  concen- 
tration of  cupric  sulphate  round  the  positive  pole  anode,  and 
(2)  a  transfer  of  copper  from  the  copper  anode  to  the  cathode 
(see.  p.  511). 


COPPER 


511 


A  copper  electrotype  of  a  medal  (or  a  page  of  type)  is  made  by 
first  preparing  a  cast  of  the  medal  in  plaster  of  Paris,  gutta  percha, 
or  wax.  The  surface  of  the  cast  is  then  rubbed  with  graphite,  to 
render  it  a  conductor,  and  the  cast  is  then  used  as  the  cathode  in  a 
cell  with  a  copper  anode,  like  that  just  described.  The  deposit  of 
copper,  when  heavy  enough,  is  stripped  off.  In  making  book 
plates,  the  cast  is  made  with  wax,  and  the  copper  electrotype  is 
strengthened  and  thickened  by  filling  the  back  with  melted  lead.* 

Copper  Refining.  —  The  tenacity,  ductility,  and  conductivity 
of  copper  are  seriously  affected  by  small  amounts  of  impurities, 
such  as  cuprous  oxide  or  sulphide,  which  are  soluble  in  the  molten 
metal.  Arsenic  amounting  to  0.03  per  cent  lowers  the  conductance 
about  14  per  cent.  There  are  also  silver  and  gold  in  smelter 
copper.  Hence,  a  large  proportion  of  the  copper  on  the  market  is 
purified  by  electrolysis.  The  principle  is  the  same  as  that  used 
in  electrotyping.  Thin  sheets  of  copper  form  the  cathodes,  and 
thick  plates  of  copper  the  anodes.  These 
are  suspended  alternately  and  close  to- 
gether in  large  troughs,  lined  with  lead, 
and  filled  with  cupric  sulphate  solution 
(Fig.  121,  diagrammatic,  view  from  above). 
The  cathodes  are  all  connected  with  the 
negative  wire  of  the  dynamo,  and  the 
anodes  with  the  positive  one.  The  Cu++ 
is  attracted  to  the  cathodes  and  is  deposited 
upon  them.  The  864—  migrates  towards 
the  anodes,  where  copper  from  the  thick 
plate  becomes  ionized  in  equivalent 
amount.  The  stock  of  cupric  sulphate 
thus  remains  the  same,  and  the  liquid  is 
stirred  to  keep  the  sulphate  from  accumu- 
lating close  to  the  anodes.  The  practical  effect  of  the  electrolysis  is 
to  carry  copper  across  from  one  plate  to  the  other.  The  cathodes 
are  removed  from  time  to  time,  and  the  deposit  of  copper  is 
stripped  from  their  surface.  Fresh  anodes  are  substituted  when 
the  old  ones  are  eaten  away.  Since  there  is  no  final  decomposi- 

*  For  newspapers,  a  plate  is  made  from  the  cast  of  the  type  more  quickly 
by  means  of  melted  stereotype  metal  (lead,  antimony,  tin;  82  :  15  :  3). 


FIG.  121. 


512  COLLEGE    CHEMISTRY 

tion  of  any  cupric  sulphate,  the  only  electrical  energy  required 
is  that  necessary  to  overcome  the  friction  of  the  moving  ions. 
Hence,  a  very  small  difference  in  potential  (less  than  0.5  volts) 
is  sufficient  (see  p.  549). 

The  less  active  metals  which  are  mixed  with  the  copper  in  the 
anode  are  not  ionized,  because  there  is  plenty  of  the  more  active 
copper  to  carry  the  current.  These  metals,  and  traces  of  sulphides, 
therefore,  fall  to  the  bottom  of  the  vat  as  a  sludge.  Zinc  and  other 
metals  more  active  than  copper,  however,  are  ionized.  Conversely, 
at  the  cathode,  the  copper,  being  the  least  active  metal  present 
in  ionic  form,  is  alone  deposited.  There  is  no  tendency  to  dis- 
charge zinc  or  hydrogen,  for  example,  so  long  as  there  are  plenty 
of  the  more  easily  discharged  copper  ions  available  (see.  p.  549). 
In  this  way,  copper,  99.8  per  cent  pure,  is  obtained,  gold  and  silver 
are  recovered  from  the  sludge,  and  the  bath  liquid  is  removed 
from  time  to  time  for  purification  from  the  more  active  metals  it 
acquires. 

SILVER 

Chemical  Relations  of  the  Element.  —  This  element  presents 
a  curious  assortment  of  chemical  properties.  It  differs  from  copper 
in  having  a  strongly  basic  oxide,  and  in  giving  salts  with  active 
acids  which  are  not  hydrolyzed  by  water.  In  these  respects  it 
approaches  the  metals  of  the  alkalies  and  alkaline  earths.  It 
resembles  copper  in  entering  into  complex  compounds,  and  in 
giving  insoluble  halides.  It  differs  from  both  copper  and  the 
metals  of  the  alkalies,  and  resembles  gold  and  platinum,  in  that 
its  oxide  is  easily  decomposed  by  heat,  with  formation  of  the  free 
metal,  and  in  the  low  position  it  occupies  in  the  electromotive 
series  and  the  consequent  slight  chemical  activity  of  the  free 
metal. 

Occurrence.  —  Native  silver,  usually  scattered  through  a  rocky 
matrix,  contains  varying  amounts  of  gold  and  copper.  Native 
copper  always  contains  dissolved  silver.  Sulphide  of  silver  (Ag2S) 
occurs  alone  and  dissolved  in  galenite  (PbS).  Smaller  amounts  of 
the  metal  are  obtained  from  pyrargyrite  Ag3SbS3,  proustite 
Ag3AsS3,  and  horn-silver  AgCl.  The  chief  supplies  come  from 
California,  Australia,  and  Mexico. 


SILVER 


513 


Metallurgy.  —  The  silver  contained  free,  or  as  sulphide,  in 
jres  of  copper  and  lead,  is  found  in  the  free  state  dissolved  in  the 
netals  extracted  from  these  ores,  and  is  secured  by  refining  them, 
[n  the  electrolytic  refining  of  copper,  silver  is  obtained  from  the 
nud  deposited  in  the  baths  (p.  512).     The  proportion  present  in 
ead  is  usually  small.     Parke's  process,  by  which   the  silver  is 
separated  from  the  lead,  takes  advantage  of  the  fact  that  molten 
zinc  and  lead  are  practically  insoluble  in  one  another,  while  silver 
s  much  more  soluble  in  zinc  than  in  lead.     Lead  dissolves  1.6  per 
cent  of  zinc,  and  zinc  1.2  per  cent  of  lead.     The  principle  is,  the 
same  as  in  the  removal  of  iodine  from  water  by  ether  (p.  129). 
The  lead  is  melted  and  thoroughly  mixed  by  machinery  with  a 
small  proportion  of  zinc.     After  a  short  time  the  zinc  floats  to  the 
top,  carrying  with  it  almost  all  of  the  silver,  and  solidifies  at  a 
temperature  at  which  the  lead  is  still  molten.     The  zinc-silver 
alloy,  largely  a  compound  Ag2Zn5,  is  skimmed  off,  and  heated 
moderately  in  a  furnace  to  permit  the  adhering  lead  to  drain  way.' 
The  zinc  is  finally  distilled  off  in  clay  retorts,  and  the  lead  remain- 
ing with  the  silver  is  removed  by  cupellation.     This  operation 
consists  in  heating  the  molten  metal  strongly  in  a  blast  of  air.     The 
lead  is  converted  into  litharge  (PbO),  which  flows  in  molten  con- 
dition over  the  edge  of  the  cupel,  and  the  silver  is  then  cast. 

Ores  of  silver  which  do  not  contain  much  or  any  lead  are  often 
smelted  with  lead  ores,  and  the  product  is  treated  as  described 
above,  but  many  other  processes  are  in  use.  The  gold,  which  goes 
with  the  silver  in  Parke's  process,  is  separated  electrolytically 
(p.  511).  Plates  of  the  silver-gold  alloy  form  the  anode,  and  silver 
nitrate  solution  the  vat-liquid.  The  silver,  being  the  more  active 
metal,  is  ionized  and  deposited  on  the  cathode,,  while  the  gold 
collects  as  a  powder  in  a  bag  surrounding  the  anode. 

During  the  first  half  of  the  nineteenth  century  the  world's  total 
output  of  silver  averaged  only  643  tons  per  year.  Up  to  1870  a 
gram  of  gold  could  buy  15.5  g.  of  silver.  Now  that  the  produc- 
tion has  reached  7800  tons,  the  same  amount  of  gold  purchases 
about  40  g.  The  chief  sources  (1911)  are  Mexico  2460  tons,  United 
States  1880,  Canada  1018,  Europe  525. 

Physical  Properties.  —  Pure  silver  is  almost  perfectly  white. 
It  melts  at  960°.  Its  sp.  gr.  is  10.5.  Its  ductility  is  such  that  wires 


514  COLLEGE    CHEMISTRY 

can  be  drawn  so  fine  that  2  kilometers  weigh  only  about  1  g.  In 
the  molten  condition  it  absorbs  mechanically  about  twenty-two 
times  its  own  volume  of  oxygen,  but  gives  up  almost  all  of  this  as 
it  solidifies.  Fantastically  irregular  masses  result  from  the 
"sprouting"  or  "spitting"  which  accompanies  the  escape  of  the 
gas. 

By  addition  of  ferrous  citrate  to  silver  nitrate,  a  red  solution  and 
lilac  precipitate  of  free  silver  can  be  made.  The  latter,  after 
washing  with  ammonium  nitrate  solution,  gives  a  red,  colloidal 
solution  in  water  (cf.  p.  416).  It  is  a  negatively  charged  colloid, 
and  is  coagulated  by  bivalent  positive  ions. 

Silver  is  alloyed  with  copper  to  render  it  harder.  The  silver 
coinage  of  the  United  States  and  the  continent  of  Europe  has  a 
"fineness  of  900"  (900  parts  of  silver  in  1000),  and  that  of  Great 
Britain  925.  Silver  ornaments  have  a  fineness  of  800  or  more. 

Chemical  Properties.  —  Silver,  when  cold,  is  oxidized  by 
ozone,  but  not  by  oxygen  (see  silver  oxide).  It  does  not  ordinarily 
displace  hydrogen  from  aqueous  solutions  of  acids.  Sulphur  com- 
pounds in  the  air  tarnish  the  surface,  producing  Ag2S,  as  do  also 
eggs,  secretions  from  the  skin  (proteins,  p.  422),  and  vulcanized 
rubber.  Silver  interacts  with  cold  nitric  acid  and  with  hot,  con- 
centrated sulphuric  acid,  giving  the  nitrate  or  sulphate  of  silver 
and  oxides  of  nitrogen  or  of  sulphur  (pp.  354,  276). 

The  Halides  of  Silver.  —  The  chloride  AgCl,  bromide  AgBr, 
and  iodide  Agl  are  formed  as  curdy  precipitates  when  a  salt  of  silver 
is  added  to  a  solution  containing  the  appropriate  halide  ion.  The 
first  is  white,  a»d  melts  at  about  457°.  The  second  and  third  are 
very  pale-yellow  and  yellow  respectively.  The  insolubility  in 
water,  which  is  very  great,  increases  in  the  above  order. 

When  exposed  to  light,  the  chloride  becomes  first  violet  (col- 
loidal silver,  dispersed  in  the  AgCl)  and  finally  brown,  chlorine 
being  liberated.  The  bromide  and  iodide  behave  similarly.  Solid 
silver  chloride  absorbs  ammonia,  forming  at  low  pressures  2 AgCl, - 
3NH3,  and  with  higher  pressures  of  ammonia  AgCl,3NH3. 

Complex  Compounds  of  Silver.  —  Silver  chloride  dissolves 
easily  in  excess  of  ammonium  hydroxide,  giving  the  complex  cation 


SILVER 


515 


Ag(NH3)2+.  The  bromide,  which  is  less  readily  soluble,  gives  the 
same  complexion.  The  iodide  is  hardly  soluble  at  all.  Ammonio- 
argentic-ion  Ag(NH3)2+,  in  solutions  of  concentrations  such  as  are 
commonly  used  (Q.1N  to  N),  gives  about  the  same  concentration  of 
argention  Ag+  as  does  the  bromide,  and  much  more  tha'a  the  highly 
insoluble  iodide  (cf.  p.  506).  Hence  the  latter  is  almost  insoluble 
in  ammonium  hydroxide,  and  can  be  precipitated  in  ammoniacal 
solution.  All  three  of  the  insoluble  halides  interact  with  solutions 
of  potassium  cyanide  and  of  sodium  thiosulphate,  and  go  into 
solution,  as  do  also  all  the  other  insoluble  silver  salts.  Usually  an 
equivalent  amount  of  the  cyanide  or  thiosulphate  suffices,  but  for 
complete  interaction  with  the  sulphide  an  excess  is  required.  With 
the  cyanide,  double  decomposition  gives  first  the  insoluble  silver 
cyanide  AgCN,  which  then  dissolves,  forming  the  soluble  potassium 
argenticyanide  K.Ag(CN)2.  The  thiosulphate  gives  a  solution 
containing  the  complex  salt  Na3.Ag(S2O3)2.  The  more  active 
metals,  like  zinc  and  copper,  displace  silver  from  all  solutions, 
whether  the  solutions -contain  simple  or  complex  salts. 

Oxides  of  Silver.  —  When  sodium  hydroxide  is  added  to  a 
solution  of  a  salt  of  silver,  a  pale-brown  precipitate  is  obtained, 
which,  after  being  freed  from  water,  is  found  to  be  argentic  oxide 
Ag2O,  and  not  AgOH.  The  aqueous  solution  of  argentic  oxide, 
however,  is  distinctly  alkaline,  and  presumably  therefore  does 
contain  the  hydroxide:  2AgOH  <=±  Ag20  +  H20.  It  is  an  active 
basic  oxide.  When  moSt,  it  absorbs  carbon  dioxide  from  the  air. 
With  ammonium  hydroxide  it  forms  the  soluble  Ag(NH4)2.OH. 
When  the  oxide  is  heated,  it  gives  off  oxygen,  leaving  metallic  silver. 
The  action  is  reversible  and  at  302°  the  dissociation  pressure  of  the 
oxygen  is  20.5  atmospheres.  At  a  higher  pressure  than  this,  there- 
fore, oxygen  will  combine  with  silver  (at  302°). 

Silver  peroxide  Ag2O2  is  formed  by  the  action  of  ozone  on  silver. 
In  the  electrolysis  of  silver  nitrate  a  deposit  of  shining  black 
crystals  which  contain  some  silver  peroxide  is  formed  on  the 
anode. 

Salts  of  Silver.  —  Silver  nitrate  AgN03  is  obtained  by  treating 
silver  with  aqueous  nitric  acid: 

3Ag  +  4HN03  ->  3AgN03  +  NO  +  2H20. 


516  COLLEGE   CHEMISTRY 

From  the  solution,  colorless  rhombic  crystals  are  deposited. 
These  melt  at  208.6°.  Thin  sticks  made  by  casting  (lunar  caustic) 
are  used  to  cauterize  sores,  because  the  substance  combines  with 
proteins  to  form  insoluble  compounds.  The  aqueous  solution  is 
neutral.  The  pure  salt  is  not  affected  by  light,  but  when  deposited 
on  cloth,  on  the  skin  of  the  fingers,  or  on  the  mouth  of  the  reagent 
bottle,  it  is  reduced  by  organic  matter,  and  silver  is  liberated.  For 
this  reason  it  is  an  ingredient  in  some  marking-inks. 

Silver  carbonate,  the  neutral  salt  Ag2CO3,  and  not  a  basic  car^ 
bonate,  is  precipitated  from  solutions  of  salts  of  silver  by  soluble 
carbonates.  It  is  slightly  yellow  in  color.  With  water  it  gives  a 
faint  alkaline  reaction  and,  like  calcium  carbonate,  is  soluble  ir 
excess  of  carbonic  acid  (p.  384).  When  heated,  the  carbonate 
decomposes,  leaving  metallic  silver.  The  sulphate  Ag2S04  is  made 
by  the  action  of  concentrated  sulphuric  acid  on  the  metal.  When 
it  is  mixed  with  a  solution  of  aluminium  sulphate  (q.v.),  octahedral 
crystals  of  silver-alum  Ag2S04,  A12(SO4)3,24H2O  are  obtained.  Silver 
sulphide  Ag2S  is  precipitated  by  hydrogen  sulphide  from  solutions 
of  all  silver  compounds,  whether  free  acids  are  present  or  not,  and 
irrespective  of  the  form  in  which  the  silver  is  combined.  Excess  of 
potassium  cyanide,  however,  prevents  its  precipitation  from  the 
argenticyanide.  The  sulphide  is  formed  by  the  action  of  metallic 
silver  on  alkaline  hydrosulphides,  and  this  interaction  forms  the 
basis  of  the  "hepar"  test  for  sulphur.  Silver  orthophosphate 
Ag3P04  (yellow),  arsenate  Ag3As04  (brown),  and  chromate  Ag-jCrC^ 
(crimson)  are  produced  by  precipitation,  and  their  distinctive 
colors  enable  us  to  use  silver  nitrate  in  analysis  as  a  reagent  for 
identifying  the  acid  radicals. 

Electroplating.  —  The  process  is  similar  to  the  electrode- 
position  of  copper  (p.  510).  The  article  to  be  plated  is  cleaned 
with  extreme  care  and  attached  to  the  negative  wire.  A  plate  of 
silver  forms  the  positive  electrode  and,  since  simple  salts  of  silver 
do  not^give  coherent  deposits,  the  bath  is  a  solution  of  potassium 
argenticyanide.  The  potassium-ion  K+  migrates  to  the  negative 
wire  and,  since  potassium  requires  a  much  greater  E.M.F.  for  its 
liberation  than  does  silver,  silver  is  there  deposited  from  the  trace 
of  argentic-ion  given  by  the  complex  silver  ions  in  the  neighbor- 
hood: Ag(CN)2-±=»Ag+  +  2CN-  Ag++0-»Ag°. 


SILVER  517 

At  the  positive  electrode  silver  goes  into  solution  in  equivalent 
amount,  giving  argentic-ion,  and  the  above  equations  are 
reversed. 

Mirrors  are  silvered  through  the  reduction  of  ammonio-silver 
nitrate  by  organic  compounds  such  as  formaldehyde  CH20  (forma- 
lin), or  grape  sugar: 

4AgOH  +  CH20  -»  3H2O  +  4Ag|  +  C02. 
The  film  of  silver  is  washed,  dried,  and  varnished. 

Photography.  —  Bromo-gelatine  dry  plates  are  covered  with 
an  emulsion  of  gelatine  in  which  silver  bromide  is  suspended. 

After  exposure,  often  for  only  a  fraction  of  a  second,  there  is  no 
visible  alteration  in  the  film.  The  image  is  developed.  Chemically, 
this  consists  in  reducing  the  silver  bromide  to  metallic  silver  by 
means  of  reducing  agents.  While  the  whole  of  the  halide  upon  the 
plate  is  reducible,  if  the  reducing  agent  is  kept  upon  it  for  a  suffi- 
cient length  of  time,  the  parts  reached  by  the  light  are  affected  first, 
and  with  a  speed  proportional  to  the  intensity  of  the  illumination 
undergone  by  each  part.  The  unreduced  silver  bromide  is  then 
dissolved  out  with  sodium  thiosulphate  ("hypo"),  and  the  silver 
image  remains.  It  is  also  thus  saved  from  being  fogged  over  by  the 
silver  that  would  be  deposited  if  the  plate  were  to  be  brought  into 
the  light  without  this  treatment  (fixing).  The  result  is  a  "  nega- 
tive/' as  the  parts  brightest  in  the  object  are  now  opaque,  and  the 
darkest  parts  of  the  object  are  transparent. 

A  common  developer  is  the  potassium  salt  of  hydroquinone 
C6H4(OH)2,  which  gives  quinone 


2AgBr  +  (KO)2C6H4  -»  2Ag  +  2KBr 

In  printing,  the  light  and  dark  are  again  reversed,  the  denser 
parts  of  the  negative  protecting  the  compounds  on  the  paper  below 
it  from  action,  and  leaving  them  white.  Either  "  bromide"  papers 
(such  as  velox,  invented  by  Baekeland),  which  require  only  brief 
exposure  and  are  developed  like  the  plate,  are  used,  or  silver 
chloride  is  the  sensitive  substance,  and  prolonged  exposure  to  light 
is  allowed  to  liberate  the  proper  amount  of  silver.  The  operation 
of  fixing  is  performed  as  before.  In  toning  chloride  papers,  a  solu- 


518  COLLEGE   CHEMISTRY 

tion  of  sodium  chloraurate  is  employed.     A  portion  of  the  silver 
dissolves,  displacing  gold  (p.  260),  which  is  deposited  in  its  place: 

NaAuCU  +  3Ag  ->  NaCl  +  3AgCl  +  Au. 
The  thin  film  of  gold  gives  a  richer  color  to  the  print. 

Analytical  Reactions  of  Silver  Compounds.  —  Argentic-ion 
Ag+  is  colorless.  Many  of  its  compounds  are  insoluble,  the  pre- 
cipitation of  the  chloride,  which  is  insoluble  in  dilute  acids,  being 
used  as  a  test.  Mercurous  chloride  and  lead  chloride  are  also  white 
and  insoluble,  but  silver  chloride  dissolves  in  ammonium  hydroxide, 
mercurous  chloride  (q.v.)  turns  black,  and  lead  chloride  is  not 
altered  in  color  (and  is  also  soluble  in  hot  water).  With  excess 
of  ammonium  hydroxide,  silver  salts  give  the  complex  cation 
Ag(NH3)2+  and,  from  solutions  containing  silver  in  this  form,  only 
the  iodide  and  sulphide  can  be  precipitated.  Sodium  thiosulphate 
and  potassium  cyanide  dissolve  all  silver  salts,  giving  salts  of 
complex  acids  with  silver  in  the  anion  (p.  515). 

GOLD  Au 

Chemical  Relations  of  the  Element.  —  This  element  forms 
two  very  incomplete  series  of  compounds  corresponding  respec- 
tively to  aurous  and  auric  oxides,  Au20  and  Au20s.  The  former  is 
a  feebly  basic  oxide,  the  latter  mainly  acid-forming.  No  simple 
salts  with  oxygen  acids  are  stable.  All  the  compounds  of  gold  are 
easily  decomposed  by  heat  with  liberation  of  the  metal.  All  other 
common  metals  displace  gold  from  solutions  of  its  compounds  (p. 
260).  Mild  reducing  agents  likewise  liberate  gold.  The  element 
enters  into  many  complex  anions. 

Occurrence  and  Metallurgy.  —  Gold  is  found  chiefly  in  the 
free  condition,  disseminated  in  veins  of  quartz,  or  mixed  with 
alluvial  sand.  Small  quantities  are  found  also  in  sulphide  ores  of 
iron,  lead,  and  copper.  Telluride  of  gold  (sylvanite),  in  which 
silver  takes  the  place  of  a  part  of  the  gold  [Au,Ag]Te2,  is  found  in 
Colorado. 

From  the  alluvial  deposits,  gold  is  usually  separated  by  washing 
in  a  cradle  (sp.  gr.,  gold  19.32,  rock  about  2.6),  as  in  the  Klondyke. 


GOLD  519 

Quartz  veins,  which  in  the  Transvaal  Colony  reach  a  thickness  of 
a  meter  and  carry  an  average  of  18  g.  of  gold  per  ton,  are  mined, 
and  the  material  is  pulverized  with  stamping  machinery.  About 
55  per  cent, of  the  gold  is  then  separated  by  allowing  the  powdered 
rock  to  be  carried  by  a  stream  of  water  over  copper  plates  amal- 
gamated with  mercury.  The  gold  dissolves  in  the  latter,  and  is 
secured  by  removal  and  distillation  of  the  amalgam.  The  45  per 
cent  of  finer  particles,  contained  in  the  sludge  which  runs  off 
("tailings"),  are  extracted  by  adding  a  dilute  solution  of  sodium 
cyanide  (MacArthur-Forest  process)  and  exposing  the  mixture  to 
the  air.  Oxidation  and  simultaneous  interaction  with  the  cyanide 
give  sodium  aurocyanide  NaAu(CN)2.  From  this  solution  the 
gold  is  isolated,  either  by  electrolysis,  or  in  the  form  of  a  purple 
powder  by  precipitation  with  zinc.  The  same  cyanide  is  used  for 
another  batch. 

The  gold  separated  from  ores  in  the  above  ways  contains  silver, 
copper,  lead,  and  other  metals,  and  various  methods  of  refining, 
mainly  electrolytic,  are  used. 

The  world's  production  of  gold  during  the  first  half  of  the  nine- 
teenth century  averaged  27  tons  annually.  In  1897  it  was  363 
tons,  and  in  1899,  472.6  tons.  It  is  partly  this  rapid  increase  in  the 
supply  of  gold  (which  is  our  standard  of  value)  which  has  made  it 
relatively  cheaper,  and  other  articles  more  expensive.  In  1913 
the  total  production  was  680  tons,  of  which  the  Transvaal  gave 
40  per  cent,  the  United  States  20  per  cent,  and  Australia  12  per 
cent. 

Properties  of  the  Metal.  —  Gold  is  yellow  in  color,  and  is  the 
most  malleable  and  ductile  of  all  the  metals.  It  melts  at  1063°. 
Its  sp.  gr.  is  19.32.  To  give  it  greater  hardness  it  is  alloyed  with 
copper,  the  proportion  of  gold  being  defined  in  "carats."  Pure 
gold  is  "24-carat."  British  sovereigns  are  22-carat  and  contain 
ft  of  copper.  American,  French,  and  German  coins  are  21.6-carat, 
or  90  per  cent  gold. 

Gold  is  not  affected  by  free  oxygen  nor  by  hydrogen  sulphide. 
It  does  not  displace  hydrogen  from  dilute  acids,  nor  does  it  interact 
with  nitric  or  sulphuric  acids  or  any  oxygen  acids  except  selenic 
acid.  It  combines,  however,  with  free  chlorine  and  bromine.  It 
interacts  with  a  mixture  of  nitric  and  hydrochloric  acids  (aqua 


520  COLLEGE    CHEMISTRY 

regia),  giving  chlorauric  acid  H.AuCl4(  =  HCl,AuCl3).  This 
happens,  not  because  aqua  regia  is  more  active  than  are  any  of  the 
substances  it  contains,  but  because  it  furnishes  both  the  chlorine 
and  the  chloride-ion  Cl~  required  to  produce  the  exceedingly  stable 
(little  dissociated)  anion  AuCLT.  Chlorine-water  (Cl2,H+,Cr,- 
C10~)  dissolves  it  also,  for  the  same  reason.  Gold  is  the  least 
active  of  the  familiar  metals. 

Compounds  with  the  Halogens.  —  Chlorauric  acid,  formed 
as  above,  is  deposited  in  yellow,  deliquescent  crystals  of 
H.AuCl4,4H20.  The  yellow  sodium  chloraurate  NaAuCl4,2H20, 
obtained  by  neutralization  of  the  acid,  is  used  in  photography  (p. 
518).  The  acid  gives  up  hydrogen  chloride  when  heated  very 
gently,  leaving  the  red,  crystalline  auric  chloride  AuCl3.  When 
dissolved  in  water,  this  gives  H2AuCl3O.  When  auric  chloride  is 
heated  to  180°,  aurous  chloride  AuCl  and  chlorine  are  formed. 
This  salt  is  a  white  powder.  It  is  insoluble  in  water,  but  in  boiling 
water  is  converted  quickly  into  auric  chloride  and  free  gold: 
3AuCl  ->  2Au  +  AuCl3  +  H2O  -»  H2AuCl30. 

Other  Compounds.  —  When  caustic  alkalies  are  added  to 
chlorauric  acid,,- or  to  sodium  chloraurate,  auric  hydroxide  Au(OH)3 
is  precipitated.  This  substance  is  an  acid,  and  interacts  with 
excess  of  the  base,  forming  aurates.  These  are  derived  from  met- 
auric  acid  (Au(OH)3  —  H20  =  HAuO2),  as,  for  example,  potassium 
aurate  K.Au02,3H20.  This  salt  interacts  by  double  decomposition 
giving,  for  instance,  with  silver  nitrate,  the  insoluble  silver  salt 
AgAu02.  Its  solution  is  alkaline  in  reaction,  showing  that  auric 
acid  is  a  weak  acid  (cf.  p.  437).  Auric  oxide  Au2O3  is  a  brown,  and 
aurous  oxide  Au2O  is  a  violet  powder.  On  account  of  its  reducing 
action,  hydrogen  sulphide  precipitates  from  chlorauric  acid  a  dark- 
brown  mixture  containing  much  aurous  sulphide  Au2S  and  free 
sulphur,  as  well  as  some  auric  sulphide  Au2S3. 

The  aurocyanides  like  KAu(CN)2(=  KCN,AuCN),  and  the 
auricyanides,  like  K.Au(CN)4(=  KCN,Au(CN)3,  are  formed  by  the 
action  of  potassium  cyanide  on  aurous  and  auric  compounds, 
respectively.  They  are  colorless  and  soluble.  Their  solutions  are 
used  as  baths,  in  conjunction  with  a  gold  anode,  for  electro- 
gilding. 


GOLD  521 

It  will  be  seen  that  gold,  although  physically  a  metal,  is  chemi- 
cally on  the  whole  a  nonmetallic  element. 

Assaying.  —  In  assaying,  the  material  containing  the  gold  is 
heated  with  borax  and  lead  in  a  small  crucible  (cupel)  of  bone  ash. 
The  lead  and  copper  are  oxidized,  and  their  oxides  are  absorbed  by 
the  cupel,  leaving  a  drop  of  molten  alloy  of  gold  and  silver.  The 
cold  button  is  flattened  by  hammering  and  rolling,  and  treated  with 
nitric  acid  to  remove  the  silver.  The  gold,  which  remains  un- 
attacked,  is  washed,  fused  again,  and  weighed.  The  acid  will  not 
interact  with  the  silver,  and  remove  it  completely,  if  the  quantity 
of  gold  exceeds  25  per  cent.  When  the  proportion  of  gold  is 
greater  than  this,  a  suitable  amount  of  pure  silver  is  fused  with  the 
alloy  (' '  quartation  ") . 

Exercises.  —  1.  Write  equations  for  the  interactions,  (a)  of  salt 
water  and  oxygen  with  copper  (p.  503),  (6)  of  ferrous  oxide  and 
sand  (p.  502). 

2.  Write  the  formulae  of  the  basic  chloride,  nitrate,  carbonate, 
and  sulphate  of  copper  as  if  these  substances  were  composed  of  the 
normal  salt,  the  oxide  and  water  (p.  369). 

3.  Can  you  develop  any  relation  between  the  racts  that  solu- 
tions of  cupric  salts  are  acid  in  reaction  and  that  they  give  basic 
carbonates  by  precipitation? 

4.  Formulate  the  action  of  potassium  cyanide  in  dissolving 
cupric  hydroxide  and  cuprous  sulphide,  assuming  that  potassium 
cuprocyanide  is  formed. 

5.  How  should  you  set  about  making  cupric  orthophosphate, 
ammonium, cuprocyanide,  and  lead  cuprocyanide? 

6.  Write  the  formulae  of  some  of  the  double  salts  analogous  to 
potassium-cupric  sulphate  (p.  509). 

7.  What  chemical  reagents  are  present  in  a  Bunsen  flame?     If 
borax  beads  were  made  in  the  oxidizing  flame  with  cupric  chloride, 
cuprous  bromide,  and  cupric  sulphate,  severally,  what  actions 
would  take  place? 

8.  Write  the  equations  for  the  interaction  of,  (a)  silver  and 
concentrated  sulphuric  acid,  (6)  silver  chloride  and  sodium  car- 
bonate when  heated  strongly,  (c)  sodium  thiosulphate  and  silver 
bromide. 


522  COLLEGE   CHEMISTRY 

x  9.  What  reagents  should  you  use  to  precipitate  the  phosphate, 
arsenate,  and  chromate  of  silver? 

10.  Write  the  equations  for  the  interactions  of,  (a)  potassium 
hydroxide  and  auric  hydroxide,  (6)  potassium  cyanide  and  sodium 
chloraurate. 

11.  In  what  respects  are  the  elements  of  this  family  distinctly 
metallic,  and  in  what  respects  are  they  allied  to  the  non-metals 
(p.  436)? 

12.  Collect  all  the  evidence  tending  to  show  that  the  cuprous 
compounds  are  more  stable  than  the  cupric. 

13.  Make  a  classified  list  of  the  methods  by  which  cupric  com- 
pounds are  transformed  into  cuprous,  and  vice  versa. 

14.  Of  which  metals  should  it  be  possible  to  obtain  colloidal 
suspensions  in  water,  and  of  which  not  (p.  260)?    Suggest  some 
liquids  in  which  you  should  expect  to  obtain  colloidal  suspensions 
of  the  alkali  metals. 


CHAPTER  XXXVIII. 

GLUCINUM,    MAGNESIUM,    ZINC,    CADMIUM,    MERCURY. 

THE   RECOGNITION   OF   CATIONS   IN   QUALITATIVE 

ANALYSIS 

The  Chemical  Relations  of  the  Family.  —  The  remaining 
elements  of  the  third  column  of  the  periodic  table,  namely  glucinum 
or  beryllium  (Gl,  or  Be,  at.  wt.  9.1),  magnesium  (Mg,  at.  wt.  24.32), 
zinc  (Zn,  at.  wt.  65.4),  cadmium  (Cd,  at  wt.  112.4),  and  mercury 
(Hg,  at.  wt.  200.6),  although  all  bivalent,  do  not  form  a  coherent 
family.  Glucinum  and  magnesium  resemble  zinc  and  cadmium, 
and  differ  from  the  calcium  family,  in  that  the  sulphates  are  soluble, 
the  hydroxides  easily  lose  water  leaving  the  oxides,  and  the  metals 
are  not  rapidly  rusted  in  the  air  and  do  not  easily  displace  hydrogen 
from  water.  They  resemble  the  calcium  family,  and  differ  from 
zinc  and  cadmium,  in  that  the  sulphides  are  hydrolyzed  by  water, 
the  oxides  are  not  reduced  by  heating  with  carbon,  complex  cations 
are  not  formed  with  ammonia,  and  the  metals  do  not  enter  into 
complex  anions.  But  glucinum  differs  from  magnesium  and 
resembles  zinc  in  that  its  hydroxide  is  acidic  as  well  as  basic.  This 
is  not  unnatural,  since  in  the  periodic  system  it  lies  between 
lithium,  a  metal,  and  boron,  a  non-metal.  Mercury  is  the  only 
member  of  the  group  that  forms  two  series  of  compounds.  These 
are  derived  from  the  oxides  HgO  and  EfeO.  Mercury  approaches 
the  noble  metals  in  the  ease  with  which  its  oxide  is  decomposed  by 
heating,  and  in  the  position  of  the  free  element  in  the  electromotive 
series. 

The  vapor  densities  of  zinc,  cadmium,  and  mercury  show  the 
vapors  of  these  three  metals  to  be  monatomic. 

GLUCINUM  Gl 

Glucinum  (or  beryllium)  is  bivalent  in  all  its  compounds.  Its 
oxide  and  hydroxide  are  basic,  and  are  also  feebly  acidic  towards 
active  bases  (see  Zinc  hydroxide).  The  element  derives  its  name 
from  the  sweet  taste  of  its  salts  (Gk.  y\vKvs,  sweet). 

523 


524  COLLEGE    CHEMISTRY 

Glucinum  occurs  in  beryl,  a  metasilicate  of  glucinum  and  alumin- 
ium Gl3Al2(Si03)6.  Beryls,  tinted  green  by  the  presence  of  a  little 
silicate  of  chromium,  are  known  as  emeralds.  The  metal,  obtained 
by  electrolysis  of  the  easily  fusible  double  fluoride  G1F2,2KF,  burns 
when  heated  in  the  air.  It  displaces  hydrogen  from  dilute  acids, 
and  when  heated,  from  caustic  potash:  G1+2KOH->K2G102+H2. 

MAGNESIUM  Mg 

Chemical  Relations  of  the  Element.  —  Magnesium  is 
bivalent  in  all  its  compounds.  The  oxide  and  hydroxide  are  basic 
exclusively.  The  element  does  not  enter  into  complex  cations  or 
anions. 

Occurrence.  —  Magnesium  carbonate  occurs  alone  as  magne- 
site,  and  in  a  double  salt  with  calcium  carbonate  MgCO3,CaC03  as 
dolomite.  The  sulphate  and  chloride  are  found  as  hydrates  and  as 
constituents  of  double  salts  (see  below)  in  the  Stassfurt  deposits. 
Olivine  is  the  orthosilicate  Mg2Si04.  Talc  (soapstone)  is  an  acid 
metasilicate  H2Mg3(Si03)4.  Serpentine  is  a  hydrated  disilicate, 
[Mg,Fe]3,Si2O7,2H2O,  as  is  also  meerschaum.  Asbestos  is  an  an- 
hydrous silicate.  The  element  derives  its  name  from  Magnesia,  a 
town  in  Asia  Minor. 

The  Metal.  —  Magnesium  is  manufactured  by  electrolysis  of 
dehydrated  and  fused  carnallite  MgCl2,KCl,6H2O.  The  iron 
crucible  in  which  the  material  is  melted  forms  the  cathode,  and  a 
rod  of  carbon  the  anode.  The  metal  is  silver-white,  and  when 
heated  can  be  pressed  into  wire  and  rolled  into  ribbon  (m.-p.  651°). 

Chemically  the  metal  is  less  active  than  are  the  metals  of  the 
alkaline  earths.  It  slowly  becomes  coated  with  a  layer  of  the  car- 
bonate. It  displaces  hydrogen  slowly  from  boiling  water  and 
rapidly  from  cold,  dilute  acids.  Magnesium  burns  in  air  with  a 
white  light.  The  ash  contains  the  nitride  Mg3N2,  as  well  as  the 
oxide. 

Powdered  magnesium  is  used  in  pyrotechny  and,  with  potassium 
chlorate  (10  :  17),  in  making  flashlight  powder  for  use  in  photog- 
raphy. 

Magnesium  Chloride  MgCl2,6HQO.  —  This  highly  deliques- 
cent salt  occurs  in  salt  deposits,  alone,  and  as  carnallite  MgCl2,- 


MAGNESIUM  525 

KC1,6H2O.  The  latter  is  an  important  source  of  potassium  chlo- 
ride (p.  445),  and  almost  all  the  magnesium  chloride  combined 
with  it  is  thrown  away.  When  the  hexahydrate  is  heated,  a  part 
of  the  chloride  is  hydrolyzed,  some  magnesium  oxide  remaining, 
and  some  hydrogen  chloride  being  given  off.  Sea-water  cannot  be 
used  in  ships'  boilers  because  of  the  hydrochloric  acid  thus  liberated 
by  the  action  of  the  magnesium  chloride  which  the  water  contains. 
Anhydrous  magnesium  chloride  MgCl2  is  obtained  by  heating  the 
double  chloride  MgCl2,NH4Cl,6H2O,  for  this  salt  can  be  dehydrated 
without  hydrolysis  of  the  chloride.  The  ammonium  chloride  is 
volatilized  (p.  453). 

The  Oxide  and  Hydroxide.  —  Magnesium  oxide  MgO  is  made 
by  heating  the  carbonate,  and  is  known  as  calcined  magnesia.  It 
is  a  white,  highly  infusible  powder,  and  is  used  for  lining  electric 
furnaces  and  making  crucibles.  It  combines  slowly  with  water  to 
form  the  hydroxide  Mg(OH)2.  . 

The  hydroxide  is  found  in  nature  as  brucite.  It  is  also  precipi- 
tated from  solutions  of  magnesium  salts  by  alkalies.  It  is  very 
slightly  soluble  in  water.  The  solution  has  a  faint  alkaline  reaction. 
When  magnesium  chloride  is  added  to  the  moist  hydroxide,  a 
hydrated  basic  chloride,  (Mg(OH)2)x,(MgCl2)y,(H20)s,  is  formed. 
The  mixture,  to  which  sawdust  is  sometimes  added,  is  used  as  a 
plaster-finish  in  building. 

Magnesium  hydroxide  is  not  precipitated  by  ammonium  hy- 
droxide when  ammonium  salts  are  present  also.  The  ammonium 
salts,  being  highly  ionized  and  giving  a  high  concentration  of 
ammonium-ion  NEU4",  repress  the  ionization  of  the  feebly  ionized 
ammonium  hydroxide,  and  so  reduce  the  concentration  of  hy- 
droxide-ion which  it  furnishes.  With  the  ordinary  concentration 
of  Mg4"1",  therefore,  the  amount  of  hydroxide-ion  existing  in  pres- 
ence of  excess  of  a  salt  of  ammonium  is  too  small  to  bring  the 
solubility  product  [Mg++]  X  [OH~]2  up  to  the  value  required  for 
precipitation  (cf.  p.  479).  Conversely,  magnesium  hydroxide 
interacts  with  solutions  of  ammonium  salts  and  passes  into  solution: 

Mg(OH)2(solid)  ±+  Mg(OH)2  (dslvd)  fc?  Mg+++2OH~ 
2NH4C1  t=?  2CP  +2NH4+ ) 

In  presence  of  excess  of  ammonium  chloride,  the  OH~  combines 
with  NHi+  to  form  molecular  ammonium  hydroxide,  and  the 


526  COLLEGE   CHEMISTRY 

equilibria  in  the  upper  line  are  displaced  forwards  to  generate  a 
further  supply  of  the  OH~.  With  sufficiently  great  concentration 
of  the  ammonium  chloride,  all  the  magnesium  hydroxide  may  thus 
dissolve.  The  whole  case  is  analogous  to  the  interaction  of  acids 
with  insoluble  salts  (p.  480). 

Other  Salts  of  Magnesium.  —  The  normal  carbonate  MgC03 
is  found  in  nature.  Only  hydrated  basic  carbonates  are  formed  by 
precipitation,  and  their  composition  varies  with  the  conditions. 
The  carbonate,  manufactured  in  large  amounts  and  sold  as  mag- 
nesia alba,  is  approximately  Mg4(OH)2(C03)3.3H20.  It  is  used  in 
medicine  and  as  a  cosmetic. 

The  common  heptahydrate  of  magnesium  sulphate  MgS04,7H2O 
crystallizes  from  cold  water  hi  rhombic  prisms,  and  is  called  Epsom 
salts.  It  is  efflorescent.  The  monohydrate  MgSO4,H2O,  which 
remains,  and  is  found  also  in  the  salt  layers  as  kieserite,  has  a  very 
low  aqueous  tension,  and  is  not  rapidly  dehydrated  except  above 
200°.  Magnesium  sulphate  is  used  in  the  manufacture  of  sodium 
and  potassium  sulphates,  and  is  employed  also  for  " loading" 
cotton  goods,  and  as  a  purgative. 

The  sulphide  MgS  may  be  formed  by  heating  the  metal  with 
sulphur.     It  is  insoluble  in  water,  but  is  decomposed  and  gives, 
finally,  hydrogen  sulphide  and  magnesium  hydroxide: 
MgS  +  2H20  <=>  Mg(OH)4  +  H2S. 

The  only  phosphate  of  importance  is  ammonium-magnesium 
orthophosphate  NH4MgP04,6H20,  which  appears  as  a  crystalline 
precipitate  when  sodium  phosphate  and  ammonium  hydroxide 
(and  chloride,  p.  525)  are  mixed  with  a  solution  of  a  magnesium 
salt. 

Analytical  Reactions  of  Magnesium  Compounds.  —  The 

magnesium  ion  is  colorless  and  bivalent.  Soluble  carbonates  pre- 
cipitate basic  carbonates  of  magnesium,  but  not  when  ammonium 
salts  are  present.  The  latter  limitation  distinguishes  compounds  of 
magnesium  from  those  of  the  calcium  family.  Sodium  hydroxide 
precipitates  the  hydroxide  of  magnesium,  except  when  salts  of 
ammonium  are  present.  The  mixed  phosphate  of  ammonium  and 
magnesium,  in  presence  of  ammonium  hydroxide,  is  the  least 
soluble  salt. 


ZINC  527 

ZINC  Zn 

Chemical  Relations  of  the  Element.  —  Zinc  is  bivalent  in  all 
its  compounds.  Of  these  there  are  two  sets,  —  the  more  numerous 
and  important  one,  in  which  zinc  is  the  positive  radical  (Zn.SC>4, 
Zn.Cl2,  etc.),  and  a  less  numerous  set,  the  zincates,  in  which  zinc 
is  in  the  negative  radical  (Na2.Zn02,  etc.).  Both  sets  of  salts  are 
hydrolyzed  by  water,  as  the  hydroxide  is  feeble  whether  it  is  con- 
sidered as  an  acid  or  as  a  base.  The  element  also  enters  into 
complex  cations  and  anions.  The  salts  are  all  poisonous. 

Occurrence  and  Extraction  from  the  Ores.  —  The  chief 
sources  of  zinc  are  calamine  Zn2Si04,H20,  smithsonite  ZnCOs, 
zinc-blende  (Ger.  blenden,  to  dazzle)  or  sphalerite  ZnS,  franklinite 
Zn(Fe02)2,  and  zincite  ZnO.  . 

The  ores  are  first  concentrated,  recently  by  froth  flotation  (p. 
503) .  They  are  then  converted  into  oxide  —  the  carbonate  by 
ignition,  and  the  sulphide  by  roasting.  The  sulphur  dioxide  is  used 
to  make  sulphuric  acid.  A  mixture  of  the  oxide  with  coal  is  then 
distilled  in  earthenware  retorts  at  1300-1400°,  the  zinc  condensing 
in  earthenware  receivers,  while  carbon  monoxide  burns  at  a  small 
opening: 

2ZnS  +  302  -> 2ZnO  +  2S02, 
ZnO  +  C     ->  CO  +  Zn. 

At  first  zinc  dust  (a  mixture  of  zinc  and  zinc  oxide)  collects  in  the 
receiver,  and  afterwards  liquid  zinc.  The  product,  which  is  cast  in 
blocks,  is  called  spelter. 

Properties  and  Uses  of  the  Metal.  —  Zinc  is  a  bluish-white 
crystalline  metal.  When  cold  it  is  brittle,  but  at  120-150°  it  can 
be  rolled  into  sheets  between  heated  rollers  and  then  retains  its 
pliability  when  cold.  At  200-300°  the  metal  becomes  once  more 
brittle,  at  419°  it  melts,  and  at  925°  it  boils.  The  vapor  at  1740° 
is  monatomic. 

The  metal  burns  in  air  with  a  greenish  flame,  giving  zinc  oxide. 
In  cold,  moist  air  it  is  very  slowly  oxidized,  and  becomes  covered 
with  a  firmly  adhering,  non-porous  layer  of  basic  carbonate  which 
protects  it  from  further  action.  The  metal  displaces  hydrogen 
from  dilute  acids.  Zinc  also  attacks  boiling  alkalies,  giving  the 
soluble  zincate  (see  below) :  2KOH  +  Zn  ->  K2Zn02  +  H2. 


528  COLLEGE    CHEMISTRY 

Sheet  zinc,  in  consequence  of  its  lightness  (sp.  gr.  7),  is  used  in 
preference  to  lead  (sp.  gr.  11.5)  for  roofs,  gutters,  and  architectural 
ornaments.  Galvanized  iron  is  made  by  dipping  sheet  iron, 
cleaned  with  sulphuric  acid  or  the  sand  blast,  into  molten  zinc. 
The  latter,  being  more  active  (p.  260),  is  rusted  instead  of  the  iron, 
but  the  rusting  is  very  slight  (see  above).  Objects  of  iron,  cleaned 
and  baked  in  zinc  dust,  also  acquire  a  coating  of  zinc  (sherardizing). 
Zinc  is  used  also  in  batteries  and  for  making  alloys  (p.  435).  It 
mixes  in  all  proportions  with  tin,  copper,  and  antimony. 

Zinc  Chloride  ZnCl2.  —  This  salt  is  usually  manufactured  by 
treating  zinc  with  excess  of  hydrochloric  acid,  evaporating  the 
solution  to  dryness,  and  fusing  the  residue.  When  hydrochloric 
acid  is  thus  present,  the  chloride  ZnCl2  is  obtained.  Evaporation 
of  the  pure  aqueous  solution,  which  is  acid  in  reaction,  results  in 
considerable  hydrolysis  and  formation  of  much  of  the  basic  chloride 
Zn2OCl2:  znci2  +  H20  <=>  HC1  +  Zn(OH)Cl,  (1) 

2Zn(OH)Cl  -»  Zn2OCl2  +  H2O.  (2) 

The  salt  is  used  in  solid  form  as  a  caustic  and,  by  injection  of  a 
solution  into  wood  (e.g.,  railway  sleepers),  as  a  poison  to  prevent 
the  growth  of  organisms  which  promote  decay.  In  both  cases  the 
salt  combines  with  proteins,  forming  solid  products.  The  hot 
solution  also  dissolves  cellulose  (cotton  or  paper).  When  the 
solution  is  pressed  through  an  orifice  into  alcohol,  the  cellulose  is 
precipitated  in  the  form  of  a  thread.  By  carbonizing  such  threads, 
carbon  filaments  for  incandescent  lamps  are  made. 

Zinc  Oxide  and  Hydroxide  and  the  Zincates.  —  The  oxide 

ZnO  is  obtained  as  a  white  powder  by  burning  zinc  or  by  heating 
the  precipitated  basic  carbonate.  It  turns  yellow  when  heated, 
recovering  its  whiteness  when  cold,  in  the  same  way  that  mercuric 
oxide  is  brown  whilst  hot  and  bright  red  when  cold.  It  is  em- 
ployed in  making  a  paint  —  zinc-white  or  Chinese  white  —  which 
is  not  darkened  by  hydrogen  sulphide.  It  is  used  also  as  a  filler  in 
making  rubber  automobile  tires. 

The  hydroxide  Zn(OH)2  appears  as  a  white,  flocculent  solid  when 
alkalies  are  added  to  solutions  of  zinc  salts.  It  interacts  as  a  basic 
hydroxide  with  acids,  giving  salts  of  zinc: 

Zn(OH)2  +  H2SO4  fc;  Zn.SO4  +  2H2O. 


ZINC  529 

It  also  interacts  with  excess  of  the  alkali  employed  to  precipitate 
it,  giving  a  soluble  zincate,  such  as  potassium  zincate  K2Zn02: 

H2Zn02T  +  2KOH  ?±  K2.Zn02  +  2H2O. 
Zinc  hydroxide  is  ionized  both  as  an  acid  and  as  a  base : 

2H+  +  Zn02=  «=fc  Zn(OH)2  (dslvd)  <±  Zn++  +  20ET 

K 

Zn(OH2)  (solid) 

Substances  which  are  both  bases  and  acids  are  called  amphoteric. 
The  ionization  as  an  acid  is  less  than  that  as  a  base,  but  both  are 
small.  Addition  of  an  acid  like  sulphuric  acid,  however,  furnishes 
hydrogen-ion;  the  hydroxyl  ions  combine  with  this  to  form  water, 
and  all  the  equilibria  are  displaced  to  the  right.  With  a  base,  on 
the  other  hand,  the  hydrogen-ion  is  removed  and  the  basic  ioniza- 
tion simultaneously  repressed,  so  that  the  equilibria  are  displaced 
to  the  left. 

Zinc  hydroxide  interacts  with  ammonium  hydroxide,  giving  the 
soluble  ammonio-zinc  hydroxide  Zn(NH3)4.(OH)2.  The  case  is  like 
those  of  copper  (p.  507)  and  silver  hydroxides  (p.  515). 

Compounds  of  zinc,  when  heated  in  the  Bunsen  flame  with  a  salt 
of  cobalt,  give  a  zincate  of  cobalt  (Rinmann's  green)  CoZn02. 

Other  Salts  of  Zinc.  —  The  normal  zinc  carbonate  ZnCO3  may 
be  precipitated  by  means  of  sodium  bicarbonate,  but  normal  car- 
bonate of  sodium  gives  basic  carbonates,  such  as  Zn2(OH)2C03: 
2ZnS04  +  2Na2C03  +  H20  ->  Zn2(OH)2C03  +  2Na2S04  +  C&T- 

Zinc  sulphate  ZnS04  is  formed  when  zinc-blende  is  roasted.  It 
gives  rhombic  crystals  of  the  hydrate  ZnSO4,7H20.  This,  and 
the  corresponding  compounds  of  magnesium  MgS04,7H20,  of  iron 
FeSO4,7H2O,  and  of  other  bivalent  metals  are  known  as  vitriols. 
The  zinc  salt  is  white  vitriol.  It  is  used  in  cotton-printing  and  as 
an  eye-wash  (f  per  cent  solution) .  The  sulphate  gives  double  salts, 
such  as  potassium-zinc  sulphate  ZnS04,K2SO4,6H2O  (c/.  p.  509). 

Zinc  sulphide  ZnS  is  more  soluble  in  water  than  is  sulphide  of 
copper,  and  hence  it  interacts  with  excess  of  strong  acids,  and 
passes  into  solution.  It  is  not  soluble  enough,  however,  to  be 
much  affected  by  weak  acids  like  acetic  acid  (cf.  p.  483).  Zinc 
sulphide  is  thus  capable  of  being  precipitated  when  acetic  acid  is 


530  COLLEGE    CHEMISTRY 

present,  or  when  hydrogen  sulphide  is  led  into  a  solution  of  the 
acetate  of  zinc: 


H2S  <=±  ZnS  J  +  2HC2H302. 

But  when  an  active  acid  is  present,  or  is  formed,  the  sulphide  is 
precipitated  incompletely  or  not  at  all,  the  action  being  reversible  : 

ZnSO4  +  H2S  <=*  ZnS  +  H2SO4. 

There  are  thus  two  ways  of  obtaining  the  sulphide  by  precipita- 
tion. A  soluble  sulphide  causes  it  to  be  thrown  down  completely, 
because  no  acid  is  liberated  in  the  action: 


ZnCl2  +  (NH4)2S  <=±  ZnS  J,  +  2NH4C1. 

The  other  method  is  to  add  sodium  acetate  to  the  solution  of  the 
salt,  and  then  lead  in  hydrogen  sulphide.  The  acid,  liberated  by 
the  action  upon  the  salt,  interacts  with  the  sodium  acetate,  giving 
a  neutral  salt  of  sodium  and  acetic  acid,  and  the  zinc  sulphide  is 
not  much  affected  by  the  latter  (cf.  p.  484).  For  uses,  see  lithopone 
(p.  497). 

Analytical  Reactions  of  Zinc  Salts.  —  Zinc  sulphide  is  pre- 
cipitated by  the  addition  of  ammonium  sulphide  to  solutions  of 
zinc  salts  and  of  zincates.  Sodium  hydroxide  gives  the  insoluble 
hydroxide,  which,  however,  interacts  with  excess  of  the  alkali, 
giving  the  soluble  zincate  of  sodium.  Compounds  of  zinc,  when 
heated  on  charcoal  with  cobalt  nitrate,  give  Rinmann's  green 
(p.  529). 

CADMIUM  Cd 

Chemical  Relations  of  the  Element.  —  This  element  is  biva- 
lent in  all  its  compounds.  Its  oxide  and  hydroxide  are  basic  exclu- 
sively, and  the  salts  are  not  hydrolyzed  by  water.  It  enters  into 
complex  compounds  having  the  ions  Cd(NH3)4++  and  Cd(CN)4=. 
Note  its  resemblances  to,  and  differences  from  zinc. 

The  Metal.  —  Aside  from  the  rare  mineral  greenockite  CdS, 
cadmium  is  found  in  small  amounts  (about  0.5  per  cent),  as  car- 
bonate and  sulphide,  in  the  corresponding  ores  of  zinc.  During 
the  reduction,  being  more  volatile  than  zinc,  it  distils  over  first 
(b.-p.  778°).  The  metal  is  white,  and  is  more  malleable  than  zinc. 


CADMIUM  531 

It  displaces  hydrogen  from  dilute  acids  (cf.  p.  260).  It  is  used  in 
making  fusible  alloys. 

Compounds  of  Cadmium. — The  chloride  CdCl2,2H20  is  efflo- 
rescent and  is  not  hydrolyzed  during  dehydration  or  in  solution. 
Zinc  chloride  (p.  528)  is  deliquescent  and  is  easily  hydrolyzed. 

The  hydroxide  Cd(OH)2  is  made  by  precipitation  (white),  and 
interacts  with  acids  (as  a  basic  hydroxide),  but  not  at  all  with 
bases.  It  dissolves  in  ammonium  hydroxide,  however,  forming 
Cd(NH3)4.(OH)2.  The  oxide  CdO  is  a  brown  powder,  obtained  by 
heating  the  hydroxide,  carbonate,  or  nitrate,  or  by  burning  the 
metal. 

The  sulphate  crystallizes  from  solution  as  3CdS04,8H20. 
Soluble  carbonates  throw  down  the  normal  carbonate  of  cadmium 
CdC03. 

Hydrogen  sulphide  precipitates  the  yellow  sulphide  CdS  even 
from  acid  solutions  of  the  salts.  The  substance  is  used  as  a  pig- 
ment. The  sulphide  of  cadmium,  however,  is  less  insoluble  in 
water  (cf.  p.  483)  than  are  the  sulphides  of  copper  and  mercury, 
and  is  not  completely  precipitated  from  a  strongly  acid  solution 
(e.g.,  HC1  >  0.32V). 

The  Solubilities  of  the  Sulphides  of  the  Metals.  —  The 

reader  will  remember  the  order  of  solubility  of  the  metallic  sul- 
phides more  easily  if  he  notes  that  it  is  practically  the  same  as  the 
order  of  activity  of  the  free  metals  (p.  260  or  Appendix  V).  •  Thus, 
the  sulphides  down  to  that  of  aluminium  are  dissolved  by  water 
(K2S  and  Na2S)  or  are  decomposed  by  water  (BaS,  SrS,  CaS,  MgS, 
A12S3).  The  hydroxides  formed,  being  soluble  (except  A1(OH3)), 
the  whole  dissolves  except  in  the  case  of  A12S3.  Zinc  sulphide  is 
insoluble  in  water,  but  is  soluble  enough  to  interact  with  (and 
dissolve  in)  dilute  acids,  even  a  feeble  one  like  acetic  acid.  Ferrous 
sulphide  requires  a  dilute  active  acid;  cadmium  sulphide  requires 
a  higher  concentration  of  an  active  acid,  as  do  also  CoS  and  NiS; 
cupric  sulphide  requires  an  oxidizing  acid  like  hot  nitric  acid;  and 
mercuric  sulphide  resists  even  this. 

i 

Analytical  Reactions  of  Cadmium  Compounds.  —  The  cad- 
mium ion  Cd"1"1"  is  bivalent  and  colorless.  The  yellow  cadmium 


532  COLLEGE    CHEMISTRY 

sulphide  is  precipitated  by  hydrogen  sulphide,  even  from  acid 
solutions  of  the  salts.  The  white,  insoluble  hydroxide  is  not 
soluble  in  sodium  hydroxide. 

MERCURY  Hg 

Chemical  Relations  of  the  Element.  —  Like  copper,  this  ele- 
ment enters  into  two  series  of  compounds,  the  mercurous  Hg1  and 
the  mercuric  Hg".  The  mercurous  halides,  like  the  cuprous  halides 
(and  the  argentic  halides),  are  insoluble  in  water  and  are  decom- 
posed by  light.  Both  of  the  oxides,  Hg2O  and  HgO,  are  basic 
exclusively,  but  in  a  feeble  degree.  The  hydroxides,  like  silver 
hydroxide,  are  not  stable,  and  lose  water,  giving  the  oxides.  The 
salts  of  both  sets  are  markedly  hydrolyzed  by  water,  and  basic 
salts  are  therefore  common.  No  carbonate  is  known.  Mercury 
enters  into  the  anions  of  a  number  of  complex  salts,  such  as 
HgCLi=,  HgI4=,  Hg(CN)4=,  etc.  It  forms  a  class  of  ammono-basic 
mercury  compounds,  like  Hg"NH2Cl,  all  of  which  are  insoluble. 

The  mercury  salts  of  volatile  acids,  like  the  corresponding  salts  of 
ammonium  (p.  345),  can  all  be  volatilized  completely.  Mercury 
vapor  and  all  mercury  compounds  are  poisonous,  the  soluble  ones 
more  markedly  so  than  the  insoluble  ones. 

Occurrence  and  Isolation  of  the  Metal.  —  Mercury  occurs 
native  and  to  a  larger  extent  as  red,  crystalline  cinnabar,  mercuric 
sulphide  HgS.  The  chief  mines  are  in  Spain,  Italy,  Austria,  and 
California. 

The  liberation  of  the  metal  is  easy,  because,  when  roasted,  the 
sulphide  is  decomposed,  and  the  sulphur  forms  sulphur  dioxide. 
The  mercury  does  not  unite  with  the  oxygen,  for  the  oxide  decom- 
poses (p.  14)  at  400-600°: 

HgS  +  02  ->  Hg  +  S02. 

In  some  places  the  ore  is  spread  on  perforated  brick  shelves  in  a 
vertical  furnace,  and  the  gases  pass  through  tortuous  flues  in  which 
the  vapor  of  the  metal  condenses. 

Physical  Properties.  —  Mercury  or  quicksilver  (N.L.  hydrargy- 
rum, from  Gk.  vd<*>p,  water,  and  apyvpos,  silver)  is  a  silver-white 
liquid.  At  -  38.7°  it  freezes,  and  at  357°  it  boils. 


MERCURY  533 

On  account  of  its  high  specific  gravity  (13.6,  at  0°)  and  low  vapor 
tension,  the  metal  is  employed  for  filling  barometers.  Its  uniform 
expansion  favors  its  use  in  thermometers.  It  forms  amalgams 
with  all  the  familiar  metals,  with  the  exception  of  iron  and  plati- 
num. The  latter,  however,  is  "wet"  by  it  (cf.  pp.  345,  519). 
Compounds,  such  as  NaHg2,  are  often  present  in  amalgams. 

Chemical  Properties.  —  When  kept  at  a  temperature  near  to 
its  boiling-point,  mercury  combines  slowly  with  oxygen.  Mercury 
does  not  displace  hydrogen  from  dilute  acids  (p.  260),  but  with 
oxidizing  acids  like  nitric  acid  and  hot  concentrated  sulphuric  acid, 
the  nitrates  and  sulphate  (mercuric)  are  formed.  With  excess  of 
mercury,  mercurous  nitrate,  and  with  excess  of  the  hot  acid,  mer- 
curic nitrate,  are  produced.  When  mercury  is  divided  into  minute 
droplets,  with  relatively  large  surface,  it  is  used  in  medicine  ("blue 
pills"),  and  shows  an  activity  which  is  entirely  wanting  in  larger 
masses. 

The  Halides  of  Mercury.  —  Mercurous  chloride  HgCl  (calomel) 
is  obtained  as  a  white  powder  by  precipitation.  It  is  made  by 
subliming  mercuric  chloride  with  mercury: 

Hg<=±2HgCl, 


.      .       .., 
or  more  usually  by  subliming  a  mixture  of  mercuric  sulphate,  made 

as  described  above,  with  mercury  and  common  salt.  It  is  de- 
posited on  the  cool  part  of  the  vessel  as  a  fibrous  crystalline  mass. 
Its  vapor  is  composed  entirely  of  mercury  and  mercuric  chloride. 
It  is  slowly  affected  by  light  just  as  is  silver  chloride.  Here,  how- 
ever, the  chlorine  which  is  released  combines  with  another  molecule 
of  the  salt  to  form  mercuric  chloride.  The  substance  is  used  in 
medicine  on  account  of  its  tendency  to  stimulate  all  organs  pro- 
ducing secretions. 

By  direct  union  with  chlorine,  mercuric  chloride  HgCl2  (corrosive 
sublimate)  is  formed.  It  is  usually  manufactured  by  subliming 
mercuric  sulphate  with  common  salt,  and  crystallizes  in  white, 
rhombic  prisms.  It  melts  at  265°  and  boils  at  307°.  The  solu- 
bility at  20°  is  7.4  :  100  Aq.  The  aqueous  solution  is  slightly  acid 
in  reaction.  The  salt  is  easily  reduced  to  mercurous  chloride. 
When  excess  of  stannous  chloride  is  added  to  the  solution,  the 


534  COLLEGE    CHEMISTRY 

white  precipitate  of  calomel,  first  formed,  passes  into  a  heavy 
gray  precipitate  of  finely  divided  mercury: 

2HgCl2  +  SnCl2  -»  SnCU  +  2HgCl, 
2HgCl  +  SnCl2 -»  SnCU  +  2Hg. 

Corrosive  sublimate,  when  taken  internally,  is  extremely  poison- 
ous. A  very  dilute  solution  (1  :  1000)  is  used  in  surgery  to  destroy 
lower  organisms  and  thus  prevent  infection  of  wounds.  Mercuric 
chloride  acts  also  as  a  preservative  of  zoological  materials,  form- 
ing insoluble  compounds  with  proteins,  and  preventing  decay. 
For  the  same  reason,  albumin  (white  of  an  egg)  is  given  as  an 
antidote  in  cases  of  sublimate  poisoning. 

Mercurous  iodide  Hgl  is  formed  by  rubbing  iodine  with  excess  of 
mercury.  It  also  appears  as  a  greenish-yellow  precipitate  when 
potassium  iodide  is  added  to  a  solution  of  a  mercurous  salt.  It 
decomposes  spontaneously  into  mercury  and  mercuric  iodide: 

2HgI  ^  Hg  +  Hgl,. 

Mercuric  iodide  HgI2  is  obtained  by  direct  union  of  mercury  with 
excess  of  iodine,  or  by  addition  of  potassium  iodide  to  a  solution  of 
a  mercuric  salt.  It  is  a  scarlet  powder,  insoluble  in  water,  but 
soluble  in  alcohol  and  ether.  It  interacts  with  excess  of  potassium 
iodide,  forming  the  soluble,  colorless  potassium  mercuri-iodide 
K2.HgLi  with  which  many  precipitants  fail  to  give  mercury  com- 
pounds. 

The  Oxides.  —  When  bases  (excepting  ammonium  hydroxide, 
see  p.  535)  are  added  to  solutions  of  mercurous  salts,  the  greenish- 
black  mercurous  oxide  Hg^O  is  thrown  down.  The  hydroxide  is 
doubtless  formed  transitorily  and  then  loses  water  (cf.  Silver  oxide, 
p.  515).  Under  the  influence  of  light  or  gentle  heat  (100°),  this 
oxide  resolves  itself  into  mercuric  oxide  and  mercury. 

Mercuric  oxide  HgO  is  formed  as  a  red,  crystalline  powder,  when 
mercury  is  heated  in  air  near  to  357°,  but  is  usually  made  by 
decomposing  the  nitrate.  Commercial  specimens,  incompletely 
decomposed,  thus  give  some  nitrogen  tetroxide  when  heated. 
It  is  formed  also  as  a  yellow  powder  by  adding  bases  (except- 
ing ammonium  hydroxide,  see  p.  535)  to  solutions  of  mercuric 
salts. 


MERCURY  535 

Other  Salts  of  Mercury.  —  Mercurous  nitrate  HgN03,H2O  is 
formed  by  the  action  of  cold,  diluted  nitric  acid  upon  excess  of 
mercury.  It  is  hydrolyzed,  slowly  by  cold,  and  rapidly  by  warm 
water,  giving  an  insoluble  basic  nitrate: 

2HgN03  +  H20  <=±  HN03  +  Hg2(OH)N03J. 

On  this  account  a  clear  solution  can  be  made  only  when  some  nitric 
acid  is  added.  Free  mercury  is  also  kept  in  the  solution  to  reduce 
mercuric  nitrate,  which  is  formed  by  atmospheric  oxidation  : 

Hg->2HgN03,     or    Hg++  +  Hg  ->  2Hg+. 


Mercuric  nitrate  Hg(N03)2,8H20  is  produced  by  using  excess  of 
warm,  concentrated  nitric  acid  with  mercury.  The  aqueous  solu- 
tion is  strongly  acid,  and  deposits  a  yellowish,  crystalline,  basic 
nitrate  Hg3(OH)20(N03)2.  The  hydrolysis  is  reversed  by  adding 
nitric  acid. 

Mercurous  sulphide  Hg2S  is  formed  by  precipitation  from  mer- 
curous  salts,  but  decomposes  into  mercury  and  mercuric  sulphide. 

Crystallized  mercuric  sulphide  HgS  occurs  as  cinnabar,  and  is 
red.  When  formed  by  precipitation  with  hydrogen  sulphide,  or  by 
rubbing  together  mercury  and  sulphur,  it  is  black  and  amorphous. 
By  sublimation,  in  the  course  of  which  it  dissociates  and  recom- 
bines,  the  black  form  gives  the  red,  crystalline  one. 

The  black  and  the  red  varieties  do  not  interact  with  concen- 
trated acids,  or  even  with  boiling  nitric  acid,  which  oxidizes  most 
sulphides  readily.  They  are,  therefore,  still  less  soluble  than  is 
cupric  sulphide  (pp.  483,  531).  They  are  attacked,  however,  by 
aqua  regia,  because  of  the  formation  of  the  negative  ion  (see  gold, 
p.  520)  of  a  complex  salt  H2.HgCl4  (=  2HCl,HgCl2).  The  red 
form  of  the  sulphide  is  used  in  making  paint  (vermilion). 

Mercuric  fulminate  Hg(ONC)2  is  obtained  as  a  white  precipitate 
when  mercury  is  treated  with  nitric  acid,  and  alcohol  is  added  to 
the  solution.  It  decomposes  suddenly  when  struck,  and  is  used  in 
making  percussion  caps  and  detonators. 

Ammono-  Compounds  of  Mercury.  —  When  ammonium 
hydroxide  is  added  to  a  solution  of  a  mercuric  salt,  a  white  sub- 
stance, of  a  type  which  we  have  not  previously  encountered,  is 
thrown  down.  Mercuric  chloride  gives  Hg(NH2)Cl,  commonly 


536  COLLEGE    CHEMISTRY 

called  " infusible  white  precipitate,"  or  ammono-basic  mercuric 
chloride. 

HgCl2  +  H.NH2  +  NH3  -»  Hg(NH2)Cl  +  NHiCl. 

The  action  is  similar  to  an  hydrolysis  which  gives  a  basic  salt: 
HgCl2  +  H.OH  -*  Hg(OH)Cl  +  HC1,  excepting  that  ammonia 
H.NH2  plays  the  part  of  the  water.  Water  gives  aquo-basic 
salts.  When  liquid  ammonia  is  the  solvent,  ammono-basic  salts 
are  produced.  In  a  few  cases,  as  here,  an  ammono-basic  salt  is 
obtained  even  when  water  is  present.  The  study  of  reactions  in 
liquid  ammonia  solutions  by  E.  C.  Franklin  has  led  to  the  discovery 
of  a  large  number  of  new  and  most  interesting  substances. 

Mercuric  nitrate  Hg(NO3)2  and  ammonium  hydroxide  give  an 
insoluble  ammono-basic  mercuric  nitrate,  Hg  =  N  — HgN03  which 
is  more  basic  than  the  foregoing: 

2Hg(N03)2  +  H3.N  +  3NH3  -»  Hg2(N)N03  +  3NH4N03. 

When  calomel  is  treated  with  ammonium  hydroxide,  it  turns  into 
a  black,  insoluble  body.  This  is  a  mixture  of  free  mercury,  to 
which  it  owes  its  dark  color,  and  "infusible  white  precipitate," 
Hg  +  Hg(NH2)Cl.  To  this  reaction  calomel  owes  its  name  (Gk. 
KaXo/xe\as,  beautiful  black).  Mercurous  nitrate  gives  a  black,  in- 
soluble mixture,  2Hg  +  Hg2(N)N03. 

Analytical  Reactions  of  Mercury  Compounds.  —  The  two 

ionic  forms  of  the  element,  mercurous-ion  Hg+  and  mercuric-ion 
Hg4^,  are  both  colorless.  Their  chemical  behavior  is  entirely 
different.  Both  give  the  black  sulphide  HgS,  which  is  insoluble  in 
acids  and  other  solvents  of  mercury  salts.  Mercurous-ion  gives 
the  insoluble,  white  chloride,  the  black  oxide,  and  a  black  mixture 
with  ammonium  hydroxide.  Mercuric-ion  gives  a  soluble  chloride, 
a  yellow,  insoluble  oxide,  and  a  white  precipitate  with  ammonium 
hydroxide.  The  behavior  with  stannous  chloride  (p.  534)  is  char- 
acteristic. With  potassium  iodide  the  two  ions  benave  differently 
(p.  534).  More  active  metals  displace  mercury  from  all  com- 
pounds. Copper  is  used  as  the  displacing  metal,  in  testing  for 
Hg+  or  Hg++,  because  the  silvery  mercury  is  easily  seen  on  its 
surface. 

Salts  of  mercury  are  volatile.  When  heated  in  a  tube  with 
sodium  carbonate,  they  give  a  sublimate  of  metallic  mercury. 


RECOGNITION   OF   CATIONS   IN   QUALITATIVE   ANALYSIS      537 

THE  RECOGNITION  OF  CATIONS  IN  QUALITATIVE  ANALYSIS 

"Wet- way"  analysis  consists  in  recognizing  the  various  positive 
and  negative  ions  present  in  a  solution  (p.  436).  In  discussing 
hydrogen  sulphide  (p.  273),  it  was  stated  that  the  sulphides  might 
be  divided  into  three  classes,  according  to  their  behavior  towards 
water  and  acids.  Now  these  differences  furnish  us  with  a  basis  for 
distinguishing  the  cations  present  in  a  solution. 

The  following  plan,  taken  in  conjunction  with  the  statements  in 
the  context,  shows  how  a  single  cation  may  be  identified,  and  how, 
when  several  cations  are  present,  a  separation  preparatory  to 
identification  may  be  effected.  What  will  be  said  applies  only  to 
the  case  of  a  solution  containing  salts  like  the  chlorides,  nitrates, 
or  sulphates  of  one  or  more  cations,  and  leaves  the  oxalates, 
phosphates,  cyanides,  and  some  other  salts,  out  of  consideration. 

Group  1.  —  Add,  first,  hydrochloric  acid,  to  find  out  whether 
cations  giving  insoluble  chlorides  are  present.  Argentic,  mer- 
curous,  and  plumbic  salts  give  the  white  AgCl,  HgCl,  and  PbCl2, 
respectively  (cf.  p.  164).  Filtration  eliminates  the  precipitate,  if 
there  is  any. 

Group  2.  —  A  free,  active  acid  being  now  present,  hydrogen 
sulphide  is  led  into  the  solution.  The  sulphides  insoluble  in  active 
acids,  namely,  HgS,  CuS,  PbS,  Bi2S3,  CdS,  As2S3,  Sb2S3,  SnS,  SnS2,  are 
therefore  thrown  down.  The  first  four  are  black  or  brown,  the  next 
two  and  the  last  are  yellow,  and  the  remaining  two  are  orange 
and  brown  respectively.  A  dark-colored  substance  will  naturally 
obscure  one  of  lighter  color,  if  more  than  one  is  present.  Filtration 
again  eliminates  the  precipitate. 

This  group  is  easily  subdivided.  Any  or  all  of  the  last  four  sul- 
phides will  pass  into  solution  when  warmed  with  yellow  ammonium 
sulphide,  for  they  give  soluble  complex  sulphides  (q.v.).  The  first 
five  sulphides,  or  any  of  them,  will  be  unaffected.  On  the  other 
hand,  these  five  sulphides,  with  the  exception  of  HgS,  will  interact 
with  hot  nitric  acid  (p.  531).  Other  reactions  are  then  used  to 
distinguish  between,  or,  if  there  is  a  mixture,  to  separate,  the  mem- 
bers of  the  sub-groups. 

Group  3.  —  The  solution  (filtrate)  is  now  neutralized  with  am- 
monium hydroxide,  and  ammonium  sulphide  is  added.  Some 
ammonium  chloride  is  also  used,  to  prevent  the  precipitation  of 


538  COLLEGE   CHEMISTRY 

magnesium  hydroxide  (p.  525),  which,  in  any  event,  would  be  in- 
complete. The  sulphides  which  are  insoluble  in  water,  and  are 
not  hydrolyzed  by  it,  now  appear.  They  are  FeS,  CoS,  NiS,  all 
black,  MnS  and  ZnS,  which  are  pink  and  white  respectively. 
There  are  precipitated  also  the  hydroxides  of  chromium  and  of 
aluminium,  Cr(OH)3  and  A1(OH)3,  because  their  sulphides  are 
hydrolyzed  by  water. 

Group  4.  —  After  filtration,  ammonium  carbonate  is  added,  and 
precipitates  the  remaining  metals  whose  carbonates  are  insoluble, 
BaCO;;,  SrCO3,  CaCO3,  with  the  exception  of  magnesium  (p.  526). 

By  addition  of  ammonium  phosphate  to  a  portion  of  the  nitrate, 
magnesium,  if  present,  now  comes  out  in  the  form  NH4MgPO4. 
There  remain  in  solution  only  salts  of  potassium,  sodium,  and 
ammonium.  Since  only  ammonium  compounds  and  other  sub- 
stances which  can  be  volatilized  have  been  added,  evaporation 
and  ignition  of  the  residue  leaves  the  salts  of  the  two  metals. 
Salts  of  ammonium  must  be  sought  in  a  fresh  sample  by  the  usual 
test  (p.  345). 

Exercises.  —  1.  Why  should  we  expect  ammonium  sulphide 
solution  to  precipitate  magnesium  hydroxide,  and  why  does  it  not 
do  so? 

2.  What  volume  of  air  is  required  to  oxidize  one  formula-weight 
of  zinc  sulphide  to  ZnO  and  SO2,  and  what  volume  of  sulphur 
dioxide  is  produced?     Is  the  gaseous  product  more  or  less  diluted 
with  nitrogen  than  when  pure  sulphur  is  burned,  and  by  how 
much? 

3.  Make  equations  showing,   (a)   the  effect  of  heating  zinc 
chloride  with  cobalt  nitrate  Co(NO3)2  in  the  Bunsen  flame  (p. 
529),  (6)  the  action  of  hydrogen  sulphide  on  sodium  zincate,  (c) 
the  actions  of  concentrated  nitric  acid  and  of  concentrated  sul- 
phuric acid  on  mercury. 

4.  What  kind  of  salts  might  take  the  place  of  sodium  acetate  in 
the  precipitation  of  zinc  sulphide  (p.  530)?     Give  examples. 

5.  Why  do  none  of  the  salts  of  the  elements  in  this  family  give 
recognizable  effects  with  the  borax  bead? 


CHAPTER  XXXIX 
ELECTROMOTIVE   CHEMISTRY 

WE  have  seen  that  many  chemical  changes  are  accompanied  by 
a  liberation  of  energy.  If  no  special  arrangement  is  made,  the 
energy  is  always  liberated  in  the  form  of  heat,  light,  and  mechan- 
ical energy.  In  changes  involving  ionogens,  however,  the  energy 
can  be  secured  in  the  form  of  electricity.  Since  the  change  sets 
an  electric  current  in  motion,  the  subject  is  called  electromotive 
chemistry.  A  knowledge  of  this  branch  of  the  science  is  essential 
for  understanding  the  numerous  commercial  applications  of 
electricity  in  chemistry.  It  also  furnishes  us  with  a  simple  method 
for  measuring  chemical  affinity  in  ionic  reactions. 

Units  of  Electrical  Energy.  —  Two  different  units  are  required 
for  denning  a  quantity  of  electrical  energy.  One  of  these  is  the 
quantity  of  electricity,  which  is  expressed  in  coulombs  (p.  237). 
The  other  is  the  electromotive  force  (E.M.F.)  of  a  current,  or  the 
difference  in  potential,  if  a  current  is  not  flowing,  or  the  flow  is  not 
being  considered.  This  is  measured  in  volts.  It  will  be  recalled 
that  in  electrolysis  equal  quantities  of  electricity  liberate  equiva- 
lent weights  of  the  component  ions  (Faraday's  law,  p.  231).  We 
shall  see,  however,  that  with  different  substances,  different  differ- 
ences in  potential  (voltages)  are  required  to  produce  the  de- 
composition. A  quantity  of  electrical  energy,  used  or  produced, 
is  expressed  by  the  product  of  the  two  factors: 

No.  of  coulombs  X  No.  of  volts  =  Quant,  of  elect,  energy  (in  Joules). 

If  we  consider  the  time  occupied  by  the  process,  the  rate  at  which 
the  electricity  flows  is  expressed  in  amperes.  One  coulomb  per 
second  is  one  ampere.  Hence: 

No.  of  amperes  X  No.  of  volts  =  Joules  per  sec.  =  Watts. 

The  kilowatt  is  1000  watts.     The  horsepower  is  736  watts. 

An  illustration  will  show  the  meaning  of  this  relation.  If  a 
50- watt  (16-candle  power)  incandescent  lamp  is  used  on  a  110- 

539 


540  COLLEGE    CHEMISTRY 

volt  circuit,  by  substituting  these  values  in  the  equation  we  per- 
ceive that  such  a  lamp  must  carry  about  0.5  amperes,  or  one 
coulomb  every  two  seconds.  If,  with  the  same  voltage,  we  wanted 
a  lamp  to  carry  more  electricity  per  second,  we  should  have  to 
reduce  the  resistance  of  the  lamp,  say,  by  shortening  the  filament, 
or  using  a  thicker  one.  Evidently,  the  number  of  such  lamps  re- 
quired to  consume  one  horsepower  would  be  736/50,  or  between 
14  and  15  lamps.  Again,  to  decompose  one  molecular  weight  of 
hydrochloric  acid  (36.5  g.)  96,540  coulombs  (p.  237)  are  required, 
and  an  E.M.F.  of  at  least  1.83  volts  (see  p.  548).  The  electrical 
energy  needed  is  therefore  96,540  X  1.83  =  176,670  joules.  If 
this  were  to  be  accomplished  by  the  current  from  a  110-volt 
direct-current  lighting  circuit,  passing  through  a  50-watt  lamp  in 
series  with  the  electrolytic  cell,  the  time  required  (x  seconds) 
would  be  given  by:  50  joules  per  sec.  X  x  sees.  =  176,670  joules, 
where  x  =  3533  seconds,  or  about  59  minutes. 

The  factors  of  electrical  energy  (volts  and  amperes)  are  easily 
measured  when  electricity  is  produced,  and  are  easily  provided 
according  to  any  specification  when  electricity  is  to  be  used. 
Hence,  it  is  much  easier  to  study  the  relations  between  chemical 
change  and  this  form  of  energy  than  between  the  same  change 
and  the  heat  or  any  other  form  of  energy  which,  under  other  con- 
ditions, it  might  produce.  Electrochemistry  is,  therefore,  in 
many  ways  better  understood,  and  easier  to  handle  than  are 
other  branches  of  chemistry  involving  energy. 

Some  Reactions  that  can  be  Used  to  Furnish  Electricity. 

—  A  few  illustrations  of  the  kinds  of  reactions  which  can  easily  be 
carried  out  in  cells,  so  as  to  furnish  an  electric  current  instead  of 
heat,  may  be  classified  thus: 

Combination  cells,  such  as  one  in  which  zinc  (or  some  other 
active  metal)  and  bromine  are  the  reacting  substances.  If  zinc 
be  placed  in  bromine- water  (or  with  pure  bromine),  we  obtain 
zinc  bromide: 

Zn  +  Br2  -» ZnBr2,     or    Zn°  +  2Br°  -»  Zn++  +  2Br~. 

Displacement  cells,  such  as  one  with  cupric  sulphate  solution 
and  a  metal  more  active  than  copper  (e.g.,  Mg,  Al,  Zn,  or  Fe), 
and  able  to  displace  (p.  260)  this  element: 

Zn  +  CuS04  ->  ZnSO4  +  Cu,     or    Zn°  +  Cu++  ->  Zn++  +  Cu°. 


ELECTROMOTIVE   CHEMISTRY 


541 


A  non-metal  may  also  be  displaced: 

2KI+Br2-»2KBr  +  I2,     or    I~  +  Br°  -»  P  +  Br". 

Oxidation  cell,  such  as  one  in  which  ferrous  chloride  FeCl2  or 
stannous  chloride  SnCl2  is  oxidized  by  chlorine-water,  giving 
FeCl3  or  SnCU: 

or    Sn++ +  201° 


Sn 


+  2C1". 


Concentration  cells,  or  cells  in  which  the  same  substance  in  two 
different  concentrations  is  used. 

The  Arrangement  of  the  Cell.  —  Every  cell  has  one  striking 
characteristic.  If  the  pairs  of  substances  mentioned  in  the  last 
section  are  placed  together,  they  interact  and  heat  is  produced. 
There  is  no  way  to  avoid  the  action,  and  the  liberation  of  the 
energy  as  heat,  if  the  substances  come  in  contact.  If,  therefore, 
all  the  energy  is  to  be  obtained  as  electrical  energy,  the  substances 
must  be  prevented  from  com- 
ing in  contact  with  one  another. 
Paradoxical  as  it  may  seem, 
it  is  easily  possible  to  obtain 
the  electricity,  and  yet  fulfill 
this  essential  condition.  The 
plan  in  all  cells  is  to  place  the 
one  substance  in  or  round  one 
pole,  and  the  other  substance 
in  or  round  the  other  pole, 
and  to  separate  the  substances 
by  a  porous  partition,  or  some 
equivalent  arrangement. 

Suppose  that  it  is  the  first 
of  the  above-mentioned  ac- 
tions that  is  to  be  used  —  the 
action  of  zinc  and  bromine. 
The  active  substances  are  ar- 
ranged as  follows:  The  pole 
on  the  left  (Fig.  122)  is  metallic  zinc.  The  solution  on  the  right 
contains  the  bromine.  The  porous  partition  in  the  center  is  per- 
meable by  migrating  ions,  but  hinders  the  mere  diffusion  of  the 


-cr+Na+- 
tl 

NaCl 


-Cr+Na+-+ 
NaCl 


Pos.  ions 


Neg.  ions*- 


FIG.  122. 


542  COLLEGE    CHEMISTRY 

dissolved  bromine  towards  the  zinc,  and  so  prevents  direct  inter- 
action with  liberation  of  heat. 

Now,  to  enable  the  cell  to  operate,  inactive,  conducting  sub- 
stances must  be  added  to  complete  the  arrangement.  A  pole  is 
added  on  the  right,  a  conducting  solution  is  placed  to  the  left  of 
the  partition,  and  a  wire  must  connect  the  two  poles.  The  wire 
may  connect  the  poles  through  a  voltmeter,  so  that  the  E.M.F. 
produced  may  be  measured.  Also,  since  bromine-water  is  a  poor 
conductor,  a  well-ionized  salt  must  be  present  along  with  the 
bromine.  The  substances  used  for  these  purposes  must  be  in- 
active. For  example,  the  pole  on  the  right  must  be  a  conductor, 
but  its  material  must  not  interact  chemically  with  the  bromine 
or  with  the  salt.  A  rod  of  carbon  or  a  platinum  wire  will  serve 
the  purpose.  A  more  active  metal,  such  as  copper,  could  not  be 
used,  because  it  would  combine  with  the  bromine.  Again,  com- 
mon salt  or  sodium  nitrate  may  be  mixed  with  the  bromine,  be- 
cause it  will  not  interact  with  bromine  or  carbon  or  platinum. 
Still  again,  the  solution  added  on  the  left  must  be  one  which  will 
not  act  upon  the  zinc  pole,  or  upon  the  solution  on  the  right, 
which  it  meets  inside  the  porous  partition.  Common  salt  or 
niter  fulfills  these  conditions.  An  acid  could  be  used  on  the  right, 
but  not  on  the  left,  for  it  would  interact  with  the  zinc.  The 
reader  should  make  a  different  selection  of  inactive  materials,  so 
as  to  become  familiar  with  the  reasoning  involved  in  the  choice 
in  each  case. 

Note  that  in  each  figure,  the  symbols  for  the  active  substances 
are  in  black-face  type,  the  products  are  in  Roman  type,  and  the 
inactive  materials  are  in  italic  type. 

The  Operation  of  the  Cell.  —  When  the  cell  has  been  as- 
sembled, and  the  wires  have  been  connected,  the  following  phe- 
nomena are  observed: 

1.  The  zinc  begins  to  form  zinc  ions,  Zn  — »  Zn4^,  an  operation 
which  leaves  the  pole  negative  (Fig.  123). 

2.  The  bromine  molecules  nearest  their  pole  touch  this  pole, 
become  bromide  ions,  Br2  — »  2Br~,  and  leave  on  the  pole  a  positive 
charge. 

3.  Since  one  pole  is  negative  and  the  other  positive,  a  current 
flows  through  the  wire. 


ELECTROMOTIVE   CHEMISTRY 


543 


u 

NaCl 


Ti 
NaCl 


4.  The  new  positive  ions  (Zn++)  round  the  left  pole  (anode) 
attract  all  the  negative  ions  in  the  cell,  and  cause  them  to  migrate 
towards  the  left  so  as  to  keep  all  parts  of  the  solution  neutral. 

5.  The   new  negative   ions 
on  the  right   (Br~)   similarly 
attract  all  the  positive  ions  in 
the  cell,  and  cause   them   to 
drift  slowly  towards  the  right 
pole  (cathode). 

6  (Very  important) .  It  will 
be  seen  that  the  zinc  and 
the  bromine  become  ionized 
at  a  distance  from  one  another 
and  do  not  actually  combine. 
The  slow  migration  of  the 
Zn++  and  Br~  ions  will,  of 
course,  after  some  hours  or 
days,  bring  some  of  these  ions 
together  in  or  near  the  parti- 
tion, and  some  molecules  will 
be  formed.  But  this  operation 
produces  no  electrical  energy 
—  it  only  gives  out  or  absorbs  heat  (p.  255).  It  is  not  an  essential 
part  of  the  operation  of  the  cell.  The  chemical  change  which  pro- 
duces the  current  is  the  ionization  of  the  two  elements,  separately. 
The  term  combination  cell  is,  therefore,  misleading.  The  cell, 
as  a  source  of  electrical  energy,  is  concerned  only  with  producing 
two  kinds  of  ions  from  the  elements.  True,  these  ions,  if  they 
united,  would  give  the  product  shown  in  the  equation  (ZnBr2), 
but  the  union,  if  it  ever  occurred,  would  be  without  electrical 
effect.  It  is  clear  that,  since  there  is  sodium  chloride  (or  some 
other  ionogen)  in  all  parts  of  the  cell,  molecules  are  ionizing,  and 
ions  are  combining,  continually,  throughout  the  whole  system. 
Thus,  on  the  left  some  zinc  chloride  molecules  are  formed  and  on 
the  right  some  sodium  bromide  molecules,  and  eventually  near  the 
center  some  zinc  bromide  molecules.  But  these  reactions  occur 
in  every  solution  containing  ionogens,  without  giving  any  current. 
In  a  cell,  the  only  reactions  which  contribute  materially  to  the 
current  are  those  taking  place  at  the  surfaces  of  the  poles. 


Pos.  ions 


Neg.  ions 


FIG.  123. 


544 


COLLEGE    CHEMISTRY 


A  Displacement  Cell.  —  In  a  similar  way,  a  cell  using  metallic 
zinc  and  cupric  sulphate  solution  may  be  arranged  (Fig.  124). 
The  zinc  forms  one  pole,  and  the  cupric  sulphate  solution  must  be 
placed  on  the  other  side  of  the  partition.  For  inactive  materials, 
a  plate  of  copper  or  of  some  metal  below  copper  in  the  activity 
series  may  be  used,  and  any  solution  (such  as  zinc  chloride  solu- 
tion) which  will  interact  neither  with  the  zinc  nor  with  the  cupric 
sulphate. 

In  following  the  operation  of  the  cell,  we  may  start  at  either 
pole.  Thus,  the  zinc  gives  zinc-ion  Zn°  — >  Zn++  -f  20 .  The  wire 


ZnCl2 


ions 


Neg. 


•*-          Pos.  Ions  -*• 


FIG.  124. 


FIG.  125. 


becomes  negatively  charged.  The  cupric-ion  is  discharged  on  the 
other  pole  GU++  — >  Cu°  +  2©,  rendering  it  positive.  All  the 
positive  ions  in  the  cell  migrate  towards  the  right  pole  (cathode). 
All  the  negative  ions  migrate  towards  the  left  pole  (anode),  since 
positive  ions  are  being  formed  on  the  left  and  are  disappearing  on 
the  right. 

When  bromine  displaces  iodine,  the  cell  may  be  arranged  as 
in  Fig.  125.  The  iodine  liberated  dissolves  in  the  potassium  iodide 
solution  and,  with  starch  emulsion  present,  its  formation  can  be 
detected  in  a  few  seconds. 


ELECTROMOTIVE    CHEMISTRY 


545 


The  Oxidation  Cell.  —  The  arrangement  whereby  stannous- 
ion  Sn++  is  oxidized  by  chlorine-water  to  stannic-ion  Sn++++  is 
shown  in  Fig.  126.  The  chlorine  01°  encountering  the  pole  be- 
comes negatively  charged,  leaving  the  pole  positive.  This  posi- 
tive charge  is  shared  by  the 
whole  conducting  wire  and, 
at  the  other  pole,  furnishes  " 
the  positive  electricity  re- 
quired to  raise  the  charge  of 
each  tin  ion  from  Sn++  to 
Sn++++.  Only  the  tin  ions 
which  touch  the  pole  can  ac- 
quire the  charge.  -Sn++ 


2C7-+Sn++ 

IT 


Na 


aCl 


Facts  Concerning  all 
Cells.  —  If  the  wire  is  discon- 
nected, the  progress  of  the 
chemical  action  is  stopped, 
although  the  difference  in  po- 
tential remains.  The  charge 
conferred  upon  a  pole,  such  "**•  lons  ~* 
as  that  from  the  cupric  ion  FlG-  126- 

(Fig.  124),  must  be  conducted  away,  before  additional  charges 
will  be  transferred  to  it. 

If  a  glass  partition  is  substituted  for  a  porous  one,  the  cell 
ceases  to  generate  electricity.  The  partition  must  permit  the 
trans-migration  of  the  ions,  which  is  a  necessary  part  of  the  oper- 
ation of  the  cell. 

When  the  circuit  is  closed,  the  changes  described  go  on  until 
one  of  the  active  materials  is  exhausted  —  for  example,  until  all 
the  cupric-ion  has  been  deposited  as  copper,  or  until  all  the  zinc 
has  been  consumed. 

The  quantity  of  electricity  producted  is  96,540  coulombs  for  each 
equivalent  weight  of  the  active  materials  transformed,  e.g.,  for 
every  65.4/2  g.  of  zinc  consumed.  The  rate  at  which  the  elec- 
tricity is  produced  is,  in  general,  greater  the  larger  the  area  of 
the  poles.  The  amperage  of  a  single  cell  is,  in  general,  very  low. 

The  E.M.F.  of  the  cell  is  not  changed  by  altering  the  size  or 
shape  of  the  poles,  or  by  using  more  or  less  of  the  solutions.  It 


546  COLLEGE   CHEMISTRY 

is  affected  by  any  change  in  the  qualities  of  the  active  materials, 
however.  Changing  the  concentration  (see  p.  551)  of  the  cupric- 
ion  (Fig.  124)  or  of  the  bromine-water  (Fig.  125)  has  an  immediate 
effect.  So  has  substituting  one  active  metal  for  another  (see  p. 
547),  as  magnesium  for  zinc  (Fig.  123).  Even  hammering  the 
metal,  thus  making  it  denser,  has  a  slight  effect. 

Single  Potential  Differences  Produced  by  the  Metals.  - 

If  we  reconsider  the  cells  described,  we  shall  see  that  there  are 
really  two  chemical  actions  in  each  cell  and  that  these  are  to  some 
extent  independent.  We  can  leave  the  zinc  (Fig.  123)  constant, 
and  change  the  concentration  of  the  bromine  or  even  substitute 
chlorine  or  iodine  for  the  latter.  On  the  other  hand,  we  can 
leave  the  bromine-water  constant,  and  exchange  the  zinc  for  some 
other  active  metal.  Thus,  the  E.M.F.  of  every  cell  is  really  the 
resultant  of  two  effects.  Now,  these  effects  can  be  measured, 
separately. 

If  we  place  zinc  in  a  solution  of  zinc  chloride,  we  find  that  there 
is  at  once  a  difference  in  potential*  bet  ween  the  metal  and  the  solu- 
tion !  The  metal  has  an  individual  tendency  to  become  ionic  — 
a  sort  of  solution  pressure  —  and  to  form  a  few  ions,  thus  making 
the  liquid  positive  and  the  metal  negative.  In  reality,  it  is  the 
tendency  of  the  atoms  of  the  metal  to  give  up  electrons  (p.  235), 
e.g.,  Zn  —  2c  =  Zn++,  which  is  being  observed.  On  the  other 
hand,  the  ions  have  a  tendency  to  deposit  themselves,  and  a  few 
may  be  deposited  (taking  up  their  electrons  and  becoming  neutral), 
rendering  the  pole  positive  and  the  solution  negative.  If  the 
former  tendency  (the  tendency  to  give  up  electrons)  is  the  stronger 
of  the  two  (the  more  active  metals),  then  a  difference  in  potential 
is  produced,  with  the  solution  positive.  If  the  latter  tendency  is 
the  stronger  (less  active  metals)  the  solution  is  observed  to  be 
negative.  Since  raising  the  concentration  of  the  metal-ions  will 
increase  the  tendency  to  deposition  and  vice  versa,  it  is  customary 
to  take  as  the  standard  solution,  for  this  purpose,  one  in  which 
the  concentration  of  the  metal-ions  is  normal  (N).  In  the  follow- 
ing table,  the  sign  preceding  the  number  is  the  charge  of  the 
solution. 


ELECTROMOTIVE    CHEMISTRY 


547 


POTENTIAL  OF  N  SOLUTIONS  IN  CONTACT  WITH  METALS 


K 

(+  2.6) 

Na 

(+  2.4) 

Ba 

(+  2.6) 

Sr 

(+3.5) 

Ca 

(+  2.4) 

Mg 

+  1.3 

Al 

+  1.0 

Mn 

+  0.8 

Zn 

+  0.5 

(E.-M.  SERIES) 

Fe(Fe++)   +0.2 
Cd  +  0.16 

Co  +  0.05 

Ni  -  0.02? 

Pb  -  0.12 

Sn(Sn++)  -0.14 
H  -  0.24 

As  -  0.53 

Cu  (Cu++)  -  0.58 


Bi  -  0.63? 

Sb  -  0.71? 
Hg(Hg+)-0.99 

Pd  -  1.03? 

Ag  -  1.04 

Pt  -  1.10? 

Au  -  1.7? 


Thus,  opposite  Mg  we  find  +1.3.  This  means  that  when  a 
piece  of  magnesium  is  placed  in  a  solution  of  a  salt  of  magnesium, 
containing  normal  concentration  of  Mg++,  the  solution  is  posi- 
tively charged  (the  metal  negatively)  and  the  difference  in  poten- 
tial is  1.3  volts.  With  silver  in  a  solution  of  a  salt  of  silver, 
containing  normal  concentration  of  silver-ion,  the  solution  is 
negative  and  the  difference  in  potential  is  —1.04  volts. 

For  a  hydrogen  pole,  a  piece  of  palladium  saturated  with  hydro- 
gen (p.  57)  is  used.  The  values  for  the  metals  which  decompose 
water  with  ease  cannot  be  observed,  and  so  calculated  values  are 
given  in  parentheses. 

Applications:  E.M.F.  of  a  Displacement  Cell.  —  For  a  cell 
in  which  one  metal  is  going  into  solution  and  another  is  being 
deposited  —  like  that  with  zinc  and 
cupric  sulphate  —  we  can  calculate  from 
the  foregoing  data  the  E.M.F.  of  the 
cell.  Zinc  in  normal  zinc-ion  solution 
makes  the  solution  +0.5.  Copper  in 
normal  cupric-ion  solution  makes  the 
solution  —0.58.  The  anodic  and  cathod- 
ic  systems  are  here  stated  as  if  they 

worked  against  one  another.     The  com-  

bined  effect  of  the  two  is  therefore 
the  difference:  +  0.5  -  (-0.58)  =  1.08 
volts. 

The  Daniell  or  gravity  cell  (Fig.  127)  represents  this  combi- 
nation. The  copper  plate  is  at  the  bottom  and  the  zinc  is  sus- 


FIG.  127. 


548  COLLEGE    CHEMISTRY 

pended  above  it.  The  cell  is  filled  with  dilute  sodium  chloride 
solution  and  crystals  of  cupric  sulphate  are  thrown  in.  So  long 
as  the  cell  is  not  disturbed,  the  heavy,  saturated  solution  of 
cupric  sulphate  remains  at  the  bottom,  so  that  no  porous  par- 
tition is  required.  The  actual  E.M.F.  of  this  cell  is  not  exactly 
that  calculated  for  normal  solutions,  because  the  cupric  sulphate 
is  in  saturated  solution,  and  the  concentration  of  the  zinc-ion 
varies,  starting  at  zero  and  increasing  as  the  cell  is  used.  It  is, 
however,  a  little  over  1  volt. 

The  Weston  Standard  Cell  contains  a  pole  of  mercury  in  a 
saturated  solution  of  mercurous  sulphate  and  cadmium  in  contact 
with  saturated  cadmium  sulphate  solution.  For  normal  solu- 
tions, the  voltage  would  be  +0.16  -  (-0.99)  =  1.15  volts.  At 
20°  it  is  1.0183  volts. 

The  Clark  Standard  Cell  contains  zinc  and  zinc  sulphate  solution 
in  place  of  the  cadmium.  With  normal  solutions  it  would  give 
+  0.5  -  (-0.99)  =  1.49  volts.  It  actually  gives  1.434  volts. 

Single    Potential    Differences  for   Non-Metallic  Ions.  — 

The  corresponding  figures  for  the  non-metals  are: 

I         -0.78        Br         -1.32        Cl         -1.59 

Hence  the  cell  with  zinc  and  bromine-water  (p.  541)  in  presence 
of  normal  concentration  of  the  respective  ions    gives   +0.5  - 
(  —  1.32)  =  1.82  volts.     Similarly,  the  cell  in  which  bromine  dis- 
places iodine  (p.  544)  gives  -0.78  -  (-1.32)  =  0.54  volts. 

Applications:     Electrolysis:     Discharging    Potentials.  — 

When  a  solution  of  a  salt,  such  as  cupric  chloride,  is  electrolyzed, 
copper  and  chlorine  are  liberated  at  the  two  poles.  Now,  when 
the  electrolysis  has  made  some  progress,  if  the  battery  is  taken 
out,  and  the  wires  are  joined,  a  current,  the  polarization  current, 
flows.  Evidently,  the  copper  and  chlorine  liberated  in  and  round 
the  electrodes  have  made  the  arrangement  into  a  copper-chlorine 
battery  cell.  Assuming  normal  concentrations,  the  E.M.F.  of 
the  polarization  current  is  —0.58  —  (  —  1.59)  =  2. 17  volts.  Now 
this  counter-current  is  in  operation  during  the  whole  electrolysis. 
To  overcome  it,  and  maintain  the  electrolysis,  evidently  an 
E.M.F.  of  at  least  2.17  volts  from  the  battery  is  required.  This 
is  called  the  discharging  potential  for  cupric  chloride. 


ELECTROMOTIVE    CHEMISTRY  549 

Applications:  Electrolytic  Refining.  —  The  electrolytic 
process  for  refining  copper  (read  p.  511)  can  now  be  more  easily 
understood.  Both  electrodes  are  made  of  copper,  and  the  solu- 
tion contains  cupric  sulphate.  There  is,  therefore,  no  difference 
in  potential  between  the  plates,  except  a  very  small  one,  due  to 
the  fact  that  one  plate  is  pure  copper  and  the  other  impure. 
Hence  a  very  slight  E.M.F.,  sufficient  to  overcome  the  difference 
just  mentioned,  and  to  overcome.  the  friction  of  the  moving  ions, 
is  all  that  is  required,  and  0.5  volts  is  sufficient. 

As  regards  the  resulting  purification,  the  anode  of  crude  copper, 
which  is  being  consumed,  contains,  besides  copper,  small  amounts 
of  less  active  metals  like  silver  and  gold,  and  of  more  active 
metals  like  zinc.  So  far  as  the  more  active  metals  are  concerned, 
the  cell  is  like  one  with  zinc  and  cupric  sulphate  (p.  544).  It 
would  run  by  itself,  without  any  outside  current,  and  would 
actually  generate  a  current.  Hence  the  active  metals  become 
ionic  easily,  and  displace  cupric-ion  from  the  solution.  The  less 
active  metals,  on  the  other  hand,  are  not  required  for  the  trans- 
ference of  the  electricity,  since  a  great  excess  of  the  more  active 
copper  is  available.  They  also  require  a  larger  E.M.F.  for  their 
ionization  than  does  copper.  Hence,  they  remain  as  metals,  and 
drop  to  the  bottom  of  the  cell  (sludge)  as  the  anode  of  crude 
copper  wears  away. 

Applications:  Couples.  —  The  fact  that  metallic  zinc  will 
displace  hydrogen-ion  from  an  acid,  or  cupric-ion  from  cupric 
sulphate  solution  can  now  be  explained.  The  more  active  metals 
are  the  ones  which  have  the  greatest  tendency  to  become  ionic. 
Each  will  deprive  the  ions  of  a  metal  below  it  in  the  list  of  their 
electric  charges: 


Zn°  +  Cu++  -»  Zn++  +  Cu°  J,  . 

Now  we  have  noted  the  facts  (pp.  55,  626)  that  contact  with  a 
platinum  wire,  or  the  presence  of  impurities  (other  metals)  in 
the  zinc,  will  hasten  its  action.  Pieces  of  two  metals  in  contact 
with  one  another  constitute  a  couple.  With  zinc  and  platinum 
in  an  acid,  a  current  is  set  up,  like  that  of  a  short  circuited  cell. 
The  zinc  becomes  negative,  the  platinum  positive,  and  the  hydro- 
gen is  liberated  upon  the  platinum.  This  facilitates  the  action 


550  COLLEGE    CHEMISTRY 

because,  when  the  platinum  is  absent,  and  the  hydrogen  gas,  in  bub- 
bles, is  liberated  on  the  surface  of  the  zinc,  this  surface  is  only 
partly  in  contact  with  the  acid  (H+),  and  so  the  liberation  of  the 
hydrogen  is  slower. 

Galvanized  iron  is  also  a  couple.  When  rain  (dilute  carbonic 
acid)  falls  upon  it,  the  zinc,  being  the  more  active  metal  (p.  528), 
is  the  anode  and  tends  to  become  ionized  (forming  the  carbonate) . 
The  iron  is  the  cathode  and  is  not  affected.  The  carbonate,  how- 
ever, forms  a  closely  adhering  coating  on  the  zinc,  and  so  but 
little  of  this  metal  is  actually  consumed,  and  the  material  is 
therefore  durable.  On  the  other  hand,  a  sheet  of  iron,  without 
the  zinc  coating,  gives  ferrous  carbonate  which  is  easily  oxidized 
to  ferric  hydroxide  (a  base  too  weak  to  give  a  carbonate).  This 
forms  a  brittle,  porous  layer  which  does  not  mechanically  protect 
the  surface  from  further  action,  and  so  the  iron  is  finally  all  oxi- 
dized. Tin-plate  (tin  on  iron,  a  couple)  is  not  attacked  so  long 
as  the  layer  of  tin  is  nowhere  broken.  But  damaged  tin-plate 
rusts  rapidly.  There,  the  iron  is  the  more  active  metal  (p.  547) 
and  forms  carbonate  and  then  hydroxide  continuously,  while 
the  tin  remains  unaffected. 

Applications:  Measurement  of  Affinity.  —  Since  equal 
quantities  of  electricity  bring  about  (or  are  brought  about  by) 
chemical  changes  in  chemically  equivalent  weights  of  material, 
it  follows  that  the  E.-M.  forces  required  (or  produced)  are  pro- 
portional to  the  chemical  affinity.  Thus  the  activities  of  the 
metals,  expressed  in  volts  (p.  547),  are  accurate  figures  for  the 
relative  affinities  of  the  metals,  so  far  at  least  as  ionic  actions  are 
concerned.  In  point  of  fact,  they  express  also  the  approximate 
affinities  of  the  metals  in  other  actions  (pp.  60,  531)  as  well.  Again, 
by  using  different  oxidizing  agents  in  place  of  the  chlorine-water 
(p.  545)  and  noting  the  differences  in  potential,  we  can  obtain 
numbers  representing  the  relative  activities  of  various  oxidizing 
agents  towards  oxidizable  ions. 

Concentration  Cells.  —  If  two  rods  of  a  metal  (e.g.,  tin)  are 
placed  together  in  the  same  solution  of  a  salt  of  the  metal  (e.g., 
stannous  chloride  SnCl2),  there  is  no  difference  in  potential,  be- 
cause the  state  of  both  poles  is  in  all  respects  the  same.  But,  if 


ELECTROMOTIVE    CHEMISTRY 


551 


SnCZ2 
cone. 


SnCl2 
dil. 


the  solution  round  one  pole  is  more  concentrated  than  that 
round  the  other,  a  difference  in  potential  is  produced  (Fig.  128). 
The  tendencies  of  the  metallic  tin  to  form  ions  are  equal,  but 
the  pressures  of  the  stannous 
ions  are  different,  and  so,  when 
the  circuit  is  closed,  stannous 
ions  are  discharged  on  the  tin 
pole  in  the  more  concentrated 
solution,  forming  long  crystals 
of  tin,  and  tin  in  equal  amount 
from  the  pole  in  the  dilute  solu- 
tion becomes  ionic. 

The  law  which  formulates 
the  relation  between  the  two 
concentrations  and  the  E.M.F. 
produced  being  known,  it  is 
possible  to  use  the  concentra- 
tion cell  for  measuring  solu- 
bilities of  insoluble  salts. 
Thus,  we  cannot  easily  meas- 
ure the  solubility  of  silver  chlo- 
ride by  the  ordinary  method 
(p.  123),  because  evaporation  of  the  solution  may  leave  a  larger 
mass  of  impurities,  derived  from  solution  of  the  glass,  than  of  dis- 
solved silver  chloride.  Hence,  we  use  two  poles  of  silver,  place  one 
in  normal  silver  nitrate  solution  and  the  other  in  saturated  silver 
chloride  solution  (with  excess  of  the  solid),  measure  the  difference 
in  potential,  and  calculate  the  ratio  of  the  concentrations  of  silver- 
ion  in  the  two  solutions.  The  absolute  value  of  that  in  the  silver 
nitrate  solution  is  known,  and  so  the  absolute  value  of  the  Ag+ 
concentration  in  the  silver  chloride  solution  can  be  found.  Since 
silver  chloride  is  a  salt  (p.  242),  it  is  very  highly  ionized  in  so 
dilute  a  solution,  and  the  molecular  concentration  of  silver-ion  is 
practically  equal  to  the  total  molecular  concentration  of  silver, 
and  therefore  of  silver  chloride  in  the  liquid. 

Exercises.  —  1.  Make  diagrams  of  the  following  cells,  choosing 
with  care  suitable  inactive  substances  to  complete  the  arrange- 
ment: (a)  chlorine-water  and  aluminium;  (6)  chlorine- water 


Pos.  ions 


Neg. 


ions- 


Fia.  128. 


552  COLLEGE    CHEMISTRY 

and  ferrous  chloride;   (c)  zinc  and  dilute  sulphuric  acid;   (e)  chlo- 
rine-water and  potassium  bromide. 

2.  Calculate  the  E.M.F.  of  each  of  the  cells  in  Ex.  1,  assuming 
normal  solutions  to  be  present. 

3.  What  will  be  the  discharging  potentials  of  solutions  of  the 
following  substances,   assuming  N  concentrations   of  the  ions: 
(a)  manganous  chloride;  (6)  hydrogen  iodide;  (c)  ferrous  bromide; 
(e)  sodium  chloride  (hydrogen  is  liberated)? 

4.  What  weight  of  zinc  must  be  ionized  every  hour  in  a  cell  in 
order  to  produce  a  current  of  5  amperes  strength?     For  how  long 
would  500  g.  of  zinc  serve  to  maintain  this  current? 

5.  In  the  zinc-bromine  cell  (p.  543),  why  is  the  zinc  pole  called 
the  anode,  although  its  charge  with  respect  to  the  platinum  is 
negative? 


CHAPTER  XL 
ALUMINIUM  AND  THE  METALS   OF  THE  EARTHS 

THE  chief  members  of  the  family  occupying  the  fourth  column  of 
the  periodic  table  are:  boron  (B,  at.  wt.  11),  aluminium  (Al,  at.  wt. 
27.1),  gallium  (Ga,  at.  wt.  70),  indium  (In,  at.  wt.  115),  thallium 
Tl,  at.  wt.  204),  all  on  the  right  side  of  the  column;  and  scandium 
(Sc,  at.  wt.  44.1),  yttrium  (Y,  at.  wt.  89),  lanthanum  (La,  at.  wt. 
139),  on  the  left  side.  These  elements  are  all  trivalent. 

The  Rare  Elements  of  this  Family.  —  The  oxide  and  hydrox- 
ide of  boron  are  acidic  (p.  431).  Those  of  aluminium  (A1(OH)3), 
gallium  (Ga(OH)s),  indium  (In(OH)3),  and  thallium  (T1O.OH)  are 
basic,  but  behave  also  as  acids  towards  strong  bases. 

Gallium  and  indium  occur  occasionally  in  zinc-blende,  and  were 
discovered  by  the  use  of  the  spectroscope.  The  former  takes  its 
name  from  the  country  (France)  in  which  the  discovery  was -made, 
and  the  latter  from  two  blue  lines  shown  by  its  spectrum. 

Thallium  is  found  in  some  specimens  of  pyrite  and  blende.  It 
was  discovered  by  Crookes,  by  means  of  the  spectroscope,  in  the 
seleniferous  deposit  from  the  flues  of  a  sulphuric  acid  factory.  It 
received  its  name  from  the  prominent  green  line  in  its  spectrum 
(Gk.  OaXXos,  a  green  twig).  It  gives  two  complete  series  of  com- 
pounds. In  those  in  which  it  is  trivalent  (thallic  salts),  it  resem- 
bles aluminium  (q.v.) .  Thus,  the  salts  of  this  series  are  more  or  less 
hydrolyzed  by  water.  Univalent  thallium  recalls  both  sodium  and 
silver.  Thallous  hydroxide  T10H  is  soluble,  and  gives  a  strongly 
alkaline  solution,  but  the  chloride  is  insoluble  in  cold  water.  The 
solutions  of  the  thallous  salts  are  neutral.  The  metal  is  displaced 
from  its  salts  by  zinc. 

Of  the  elements  on  the  left  side  of  the  column,  scandium,  whose 
existence  and  properties  were  predicted  by  Mendelejeff  (p.  301),  is 
the  best  known.  The  metals  of  the  rare  earths,  of  which  it  is  one, 
are  found  in  rare  minerals  such  as  euxenite,  gadolinite,  orthite,  and 

553 


554  COLLEGE    CHEMISTRY 

monazite,  which  occur  in  Sweden,  Greenland,  and  the  United 
States.  Cerium  (Ce,  at.  wt.  140.25),  neodymium  (Nd,  at.  wt. 
144.3),  and  praseodymium  (Pr,  at.  wt.  140.9)  occur  along  with 
lanthanum  in  cerite,  a  silicate  of  these  four  elements.  These  four 
are  included  amongst  the  metals  of  the  rare  earths.  The  com- 
pounds of  many  of  these  rare  elements  behave  so  much  alike  that 
separation  is  difficult.  There  are  several  with  atomic  weights 
near  to  that  of  lanthanum  for  which  accommodation  cannot  easily 
be  found  in  the  periodic  table.  Ostwald  has  compared  them  to  a 
group  of  minor  planets  such  as  in  the  solar  system  takes  the  place 
of  one  large  planet. 

ALUMINIUM 

The  Chemical  Relations  of  the  Element.  —  Aluminium  is 
trivalent  exclusively.  Its  hydroxide,  like  that  of  zinc  (p.  529),  is 
amphoteric,  that  is  to  say,  it  is  feebly  acidic  as  well  as  basic,  and 
hence  the  metal  forms  two  sets  of  compounds  of  the  types  Na3.AlO3 
(sodium  aluminate)  and  A12.  (804)3.  The  salts  of  both  series  are 
more  or  less  hydrolyzed  by  water,  the  former  very  conspicuously 
so.  It  is  worth  noting  that  the  hydroxides  of  the  trivalent  metals, 
or  metals  in  the  trivalent  condition,  such  as  A1(OH)3,  Cr(OH)3, 
Fe(OH)3,  are  all  distinctly  less  basic  than  are  those  of  the  bivalent 
metals  such  as  Zn(OH)2,  Cd(OH)2,  Fe(OH)2,  Mn(OH)2.  Alumin- 
ium does  not  enter  into  complex  anions  or  cations,  and  is  too  feebly 
base-forming  to  give  salts  like  the  carbonate  or  sulphide. 

Occurrence.  —  Aluminium  is  found  very  plentifully  in  combi- 
nation, coming  next  to  oxygen  and  silicon  in  this  respect.  The 
felspars  (such  as  KAlSi308),  the  micas  (such  as  KAlSiO4),  and 
kaolin  (clay)  H2Al2(Si04)2,H20  are  the  commonest  minerals  con- 
taining it.  Since  the  soil  has  been  formed  largely  by  the  weather- 
ing of  minerals  like  the  felspars,  clay  and  other  products  of  the 
decomposition  of  such  minerals  constitute  a  large  part  of  it.  Cryo- 
lite is  a  double  fluoride  3NaF,AlF3.  Various  forms  of  the  oxide 
and  hydroxide  are  also  found. 

Preparation  and  Physical  Properties.  —  The  metal  is  now 
made  on  a  large  scale  by  electrolysis  of  the  oxide  A12O3  dissolved 
in  a  bath  of  molten  cryolite  (m.-p.  1000°),  a  process  invented  by 
C.  M.  Hall  (1886).  The  operation  is  conducted  in  ceils  (5x3  feet, 


ALUMINIUM  555 

or  larger),  the  carbon  linings  of  which  form  the  cathodes  (Fig.  129). 
The  anodes  are  rods  of  carbon  which  combine  with  the  oxygen  as 
it  is  liberated.  The  molten  metal  (m.-p.  659°)  sinks  to  the  bottom 

of  the  cell  and  is  drawn  off  periodically,     +  ____ 

while  fresh  portions  of  the  oxide  are 
added  from  time  to  time.  The  oxide  is 
made  from  bauxite  (see  below),  and 
must  be  free  from  oxide  of  iron  and 
other  impurities,  as  the  metal  cannot 
be  purified  commercially.  The  current 
(E.M.F.  5-6  volts)  maintains  the  tem- 
perature of  the  molten  materials,  and 
causes  the  decomposition.  In  1866  FlG-  129> 

aluminium  cost  $250-750  (£50-150)  per  kilogram.  In  1883  the 
whole  production  was  about  40  kilos.  In  1913  the  United  States 
alone  consumed  35  million  kilos,  costing  about  50  cents  (2/-)  per  kilo. 
The  metal  melts  at  658.5°,  but  is  not  mobile  enough  to  make 
castings.  It  is  exceedingly  light  (sp.  gr.  2.6),  and  in  tensile  strength 
excels  the  other  metals,  with  the  exception  of  iron  and  copper. 
It  is  malleable,  and  the  foil  is  taking  the  place  of  tin  foil  for  wrap- 
ping foods.  It  has  a  silvery  luster,  and  tarnishes  very  slightly, 
the  firmly  adhering  film  of  oxide  first  formed  protecting  its  surface. 
Although,  comparing  cross-sections,  it  is  not  so  good  a  conductor 
of  electricity  as  is  copper,  yet  weight  for  weight  it  conducts  better. 
It  is  difficult  to  work  on  the  lathe  or  to  polish,  because  it  sticks  to 
the  tools,  but  the  alloy  with  magnesium  (about  2  per  cent)  called 
magnalium  has  admirable  qualities  in  these  respects.  Aluminium 
bronze  (5-12  per  cent  aluminium  with  copper)  is  easily  fusible,  has 
a  magnificent  golden  luster,  and  possesses  mechanical  and  chemical 
resistance  exceeding  that  of  any  other  bronze.  The  metal  and  its 
alloys  are  used  for  making  cameras,  opera-glasses,  cooking  utensils, 
and  other  articles  requiring  lightness. and  strength,  as  well  as  for 
the  transmission  of  electricity.  The  powdered  metal,  mixed  with 
oil,  is  used  in  making  a  silvery  paint. 

/%  1  ^^  ->  '    *\_ 

Chemical  Properties.  —  The  metal  displaces  hydrogen  from 

hydrochloric  acid  very  easily.  It  displaces  hydrogen  also  from 
boiling  solutions  of  the  alkalies,  forming  aluminates: 

2A1  +  6NaOH  ->  2Na3A103  +  3H2. 


556  COLLEGE   CHEMISTRY 

In  consequence  of  its  very  great  affinity  for  oxygen,  aluminium 
displaces  from  their  oxides  the  metals  below  magnesium  in  the 
E.-M.  series.  Thus,  when  a  mixture  (thermite)  of  aluminium 
powder  and  ferric  oxide  is  placed  in  a  crucible  and  ignited  by  means 
of  a  piece  of  burning  magnesium  ribbon,  aluminium  oxide  and  iron 
are  formed : 

Fe203  +  2A1  -» A1203  +  2Fe. 

The  very  high  temperature  (about  3000°)  produced  by  the  action  is 
sufficient  to  melt  both  the  iron  (m.-p.  1530°)  and  the  oxide  of 
aluminium  (m.-p.  2050°).  The  products,  not  being  miscible, 
separate  into  two  layers.  This  very  simple  method  of  making  pure 
specimens  of  metals  like  chromium,  uranium,  and  manganese,  whose 
oxides  are  otherwise  hard  to  reduce,  is  called  by  Goldschmidt, 
the  inventor,  aluminothermy.  By  preheating  the  ends  of  steel  rails 
with  a  gasoline  torch,  firing  a  mass  of  thermite  in  a  crucible  above 
the  joint,  and  allowing  the  iron  to  flow  into  the  joint,  perfect  welds 
are  made.  In  the  same  way,  large  castings,  like  propeller  shafts, 
when  broken,  can  be  mended.  The  sulphides,  such  as  pyrite,  are 
reduced  with  like  vigor  by  aluminium. 

The  largest  part  of  the  aluminium  of  commerce  is  used  by  steel- 
makers. When  added  in  small  amount  (less  than  1  :  1000)  to 
molten  steel,  it  combines  with  the  gases,  and  gives  sound  ingots 
free  from  blow  holes. 

Aluminium  Chloride  AIC13.  —  If  the  metal  or  the  hydroxide 
is  treated  with  hydrochloric  acid,  and  the  solution  is  allowed  to 
evaporate,  the  hydrated  chloride  A1C13,6H20  is  formed.  When 
heated,  this  hydrate  is  completely  hydrolyzed,  hydrochloric  acid 
is  given  off,  and  only  the  oxide  remains.  The  anhydrous  chloride 
Aids  is  made  by  passing  dry  chlorine  over  aluminium.  Since  it 
sublimes  as  a  white  crystalline  solid  without  melting,  when  thus 
prepared  it  is  vaporized  and  condenses  in  a  cool  part  of  the  tube. 
It  fumes  when  exposed  to  moist  air  on  account  of  the  hydrogen 
chloride  produced  by  hydrolysis,  and  only  with  excess  of  hydro- 
chloric acid  does  it  give  a  clear  solution  free  from  basic  salts. 

Aluminium  Hydroxide  and  the  Aluminates.  —  When  an 
alkali  is  added  to  a  solution  of  a  salt  of  aluminium,  the  hydroxide 
A1(OH)3  is  precipitated  in  gelatinous  form.  It  loses  water  gradu- 


ALUMINIUM  557 

ally  when  dried,  forming  no  intermediate  hydroxides,  until  Al20s 
remains.  Natural  forms  of  this  substance  are  hydrargyllite 
A1(OH)8  (=  A1203,3H20),  bauxite  A120(OH)4  (=  A1203,2H20), 
which  always  contains  ferric  oxide,  and  diaspore  A10.0H  (  =  Al203r 
H20). 

Commercially,  the  hydroxide  is  made  by  heating  bauxite  with 
dry  sodium  carbonate,  and  extracting  the  sodium  metaluminate 
with  water: 

A120(OH)4  +  Na2C03  -*  2NaAlO2  +  C02  +  2H20. 

The  iron,  present  as  an  impurity,  remains,  as  ferric  oxide,  undis- 
solved.  The  hydroxide  is  then  precipitated  by  passing  carbon 
dioxide  through  the  solution: 

2NaA102  +  CO2  +  3H20  -»Na2C03  +  2A1(OH)3. 

Aluminium  hydroxide,  being  amphoteric,  interacts  both  with 
acids  and  with  bases,  and  is,  therefore,  like  zinc  hydroxide  (p.  531), 
ionized  both  as  a  base  and  as  an  acid.  It  interacts  only  slightly 
with  ammonium  hydroxide,  because  this  substance  is  too  feebly 
basic,  but,  from  the  solution  in  the  active  alkalies,  the  aluminates 
Na3.AlO3,  Na.AlO2,  and  K.A1O2,  can  be  obtained  in  solid  form. 
The  aluminates  are  largely  hydrolyzed  by  water: 

NaA102  +  2H20  +±  NaOH  +  A1(OH)3. 

When  calcium  chloride  is  added  to  a  solution  of  sodium  alumi- 
nate,  the  insoluble  calcium  metaluminate  is  deposited: 

2NaA102  +  CaCl2  -»  Ca(A102)2  +  2NaCl. 

The  relations  of  these  various  substances  are  shown  by  the  follow- 
ing formulae: 

3-H          X0-Na  0  0  <Q 

Al-O-H      Al-O-Na      Alf  Alf  )Ca 


A  number  of  insoluble  metaluminates,  such  as  spinelle  Mg(A102)2, 
and  gahnite  Zn(A102)2,  are  found  in  nature.  They  contain  bi- 
valent metals  in  place  of  the  calcium  in  the  last-named  compound. 


558  COLLEGE    CHEMISTRY 

Aluminium  Oxide  A12O3.  —  The  oxide  (alumina)  is  manu- 
factured by  heating  the  pure  hydroxide  made  from  bauxite  (see 
above).  It  is  found  in  nature  in  pure  form  as  corundum.  This 
mineral  is  only  one  degree  less  hard  than  the  diamond.  Emery 
is  a  common  variety,  contaminated  with  ferric  oxide,  and  was 
widely  used  as  an  abrasive,  until  largely  displaced  by  carborundum. 
The  ruby  is  pure  aluminium  oxide  tinted  by  a  trace  of  a  compound 
of  chromium,  while  the  sapphire  is  the  same  material  colored  with 
aluminates  of  iron  and  titanium.  It  is  said,  however,  that  the 
same  tint  is  conferred  upon  colorless  corundum  by  exposure  to  the 
influence  of  salts  of  radium.  By  ingenious  methods  of  fusing  the 
oxide,  "synthetic"  sapphires  and  rubies  are  now  made  in  large 
quantities.  Alundum,  a  refractory  material  for  crucibles,  is  made 
by  heating  objects  made  of  the  oxide  in  the  electric  furnace  until 
a  small  proportion  of  the  material  is  melted. 

Aluminium  Sulphate:  The  Alums.  —  Aluminium  sulphate 
A12(S04)3,18H20  is  prepared  by  treating  either  bauxite,  or  the  pure 
hydroxide  made  from  bauxite,  or  pure  clay  (kaolin)  with  sulphuric 
acid.  In  the  latter  case  the  insoluble  residue  of  silicic  acid  is 
removed  by  filtration: 

H2Al2(Si04)2  +  3H2S04  -»  A12(S04)3  +  2H2Si03  +  2H20. 

The  solution  of  the  sulphate  is  acid  in  reaction.  It  crystallizes  in 
leaflets  which,  when  the  source  was  clay  or  bauxite,  have  a  yellow 
tinge  due  to  the  presence  of  iron  as  an  impurity.  The  salt  is  used 
as  a  source  of  precipitated  aluminium  hydroxide  in  paper-making, 
water  purification,  and  dyeing. 

When  sulphate  of  potassium  solution  is  added  to  a  strong  solu- 
tion of  aluminium  sulphate,  octahedral  crystals  of  potash  alum 
(see  below)  are  deposited.  This  is  a  doub!6  salt,  and  is  one  of  a 
large  class  known  as  the  alums.  The  alums  have  the  general 
formula  M2IS04,M2III(S04)3,24H2O,  and  may  be  made  as  above  by 
using  a  sulphate  of  a  univalent  metal  with  one  of  a  trivalent  metal. 
Thus,  for  M1  we  may  use  K,  NH4,  Rb,  Cs,  and  Tl1,  and  for 
M111,  Al,  Fem,  Crm,  Mnm,  and  Tl111.  All  the  alums  crystallize 
in  octahedra. 

Potassium-aluminium  sulphate  K2S04,A12(S04)3,24H20,  ordinary 
alum,  is  made  from  aluminium  sulphate.  It  is  also  prepared  by 


ALUMINIUM  559 

heating  alunite,  a  basic  alum  found  near  Rome  and  in  Hungary, 
and  extracting  the  product  with  water.  The  alunite  KA13(OH)6- 
(864)2  leaves  an  insoluble  residue  of  the  hydroxide,  mixed  with 
ferric  oxide  which  is  present  as  an  impurity: 

2KA13(OH)6(S04)2  ->  K2S04,A12(S04)3  +  4A1(OH) 


3. 


The  hydrated  salt  melts  at  90°.  An  aqueous  solution  of  this 
salt,  or  of  sodium  phosphate  (p.  464),  is  used  for  fireproofing 
draperies.  The  crystals  deposited  in  the  fabric  melt  easily,  and 
the  fused  material  protects  the  fibers  from  access  of  oxygen.  When 
heated  more  strongly  alum  loses  its  water  of  hydration,  together 
with  some  sulphur  trioxide,  and  leaves  a  slightly  basic,  anhydrous 
salt  known  as  burnt  alum.  Potash-alum  and  ammonium-alum  are 
more  easily  freed  from  impurities  (e.g.,  compounds  of  iron)  by 
recrystallization  than  is  aluminium  sulphate,  and  the  alums  are 
therefore  used  instead  of  the  latter  in  medicine,  in  dyeing  delicate 
shades,  and  to  replace  cream  of  tartar  in  baking  powder  (p.  463). 
In  the  last  case,  the  reaction 

K2S04,A12(SO4)3,24H20  +  6NaHC03  -»  K2S04  +  3Na2S04 
+  2A1(OH)3  +  6C02  +  24H20 

liberates  carbon  dioxide  by  hydrolysis  of  the  aluminium  carbonate. 

Hydrolysis  of  Aluminium  Carbonate.  —  The  foregoing 
action,  and  others  discussed  above  (p.  557),  show  that  the  carbon- 
ate is  completely  hydrolyzed: 

A12(C03)3  +  6H20  ?±  2A1(OH)3  j,  +  3H2C03  ->  3H20  +  3CO2  1  . 

It  will  be  seen  that  this  may  be  due  only  in  part  to  the  feebly  basic 
qualities  of  the  hydroxide.  If  the  hydroxide  were  not  precipitated, 
it  would  cause  some  reversal  of  the  action,  and  some  of  the  car- 
bonate would  remain.  The  insolubility  of  one  product  explains 
also  other  cases  of  the  complete  hydrolysis  of  salts  (e.g.,  ammonium 
silicate  p.  429  and  next  section). 

Aluminium  Sulphide  A12SS.  —  This  salt  is  most  easily 
obtained  by  mixing  pyrite  with  aluminium  powder  and  igniting 
with  magnesium  ribbon  (p.  556)  : 

3FeS2  +  4A1  ->  2A12S3  +  3Fe. 


560  COLLEGE    CHEMISTRY 

It  forms  a  grayish-black  solid,  and  is  decomposed  by  water,  as  is 
magnesium  sulphide,  giving  the  hydroxide  and  hydrogen  sulphide. 
On  this  account,  the  sulphide,  like  magnesium  sulphide  (p.  538), 
cannot  be  formed  by  precipitation  in  presence  of  water.  Thus, 
ammonium  sulphide  with  a  salt  of  aluminium,  in  solution,  gives  a 
precipitate  of  aluminium  hydroxide: 

A12(S04)3  +  3(NH4)2S  +  3H20  ->  2A1(OH)3  j  +  3(NH4)2S04  +  3H2S. 


Coagulation  Method  of  Purifying  Water:  Sizing  Paper.  — 

When  aluminium  hydroxide  is  formed,  it  gives  first  a  colloidal  sus- 
pension, which  coagulates  to  a  gelatinous  precipitate.  When  this 
precipitate  is  produced  in  water  used  for  domestic  purposes,  and 
containing  fine,  suspended  matter,  the  gelatinous  material  causes 
the  fine  particles  to  collect  into  larger  ones  which  settle  rapidly, 
and  permits  the  use  of  relatively  small  settling  ponds.  These 
larger  particles  enclose  also  practically  all  the  bacteria.  If  the 
water  is  slightly  hard,  crude  aluminium  sulphate,  alone,  is  added: 

3Ca(HC03)2  +  A12(S04)3  ->  3CaS04  +  2A1(HC03)3,         (1) 
A1(HC03)3  +  3H20  ->  Al(OH),  j  +  3H2C03.  (2) 

If  the  water  is  very  soft,  a  little  lime  Ca(OH)2  is  added.  The  few 
remaining  bacteria  are  destroyed  by  later  addition  of  bleaching 
powder  or  chlorine-  water  (p.  312). 

In  some  localities  crude  ferrous  sulphate,  obtained  as  a  by- 
product in  cleaning  iron,  is  cheaper,  and  is  employed  instead.  The 
lime  precipitates  ferrous  hydroxide  Fe(OH)2.  This  is  quickly 
oxidized  to  colloidal  ferric  hydroxide  Fe(OH)3,  which  coagulates 
the  suspended  matter. 

Aluminium  hydroxide  is  employed  also  in  sizing  cheaper  grades 
of  paper  (p.  402),  an  operation  required  to  prevent  the  absorption 
and  consequent  spreading  of  the  ink.  For  writing-paper,  gelatine 
solution  is  employed.  In  making  printing-papers,  rosin  soap 
(made  by  dissolving  rosin  in  caustic  soda)  is  mixed  with  the  paper- 
pulp,  and  aluminium  sulphate  is  added.  The  rosin  and  aluminium 
hydroxide  are  precipitated  in  the  pulp,  perhaps  in  feeble  combina- 
tion, and  pressing  between  hot  rollers  afterwards  melts  the  former 
and  gives  a  surface  to  the  paper. 

Delicate  cloth  goods  are  rendered  waterproof  by  saturating 
them  with  aluminium  acetate  solution  and  then  steaming  them  to 


ALUMINIUM  561 

promote  hydrolysis.  The  aluminium  hydroxide  is  thus  precipi- 
tated in  the  capillaries  of  the  cotton  or  linen  and  renders  them 
non-absorbent : 

A1(C02CH3)3  +  3H20  <=±  A1(OH)8  +  3HC02CH3. 

Kaolin  and  Clay:    Earthenware  and  Porcelain.p**+Ry  the 

action  of  water  and  carbon  dioxide  upon  granite,  and  other  rocks 
containing  felspar  KAlSi3O8,  the  potash  is  slowly  removed,  and 
the  compound  is  changed  largely  into  a  hydrated  orthosilicate 
H2Al2(SiO4)2,H20.  When  pure,  it  forms  kaolin  or  china  clay,  a 
white,  crumbly  material.  When  washed  away  and  redeposited,  it 
usually  acquires  compounds  of  iron,  the  carbonates  of  calcium  and 
magnesium,  and  sand  (silica),  becoming  common  clay.  Ocher, 
umber,  and  sienna  are  clays  colored  with  oxiles  of  iron  and  man- 
ganese. Fuller's  earth  is  a  purer  variety. 

The  plasticity  of  clay,  a  property  connected  with  the  colloidal 
nature  of  the  kaolin,  enables  it  to  be  fashioned  into  various  shapes. 
When  heated,  it  shrinks  and  becomes  a  hard,  solid,  porous  mass, 
and  does  not  melt.  These  two  properties  enable  us  to  use  it  in 
making  bricks,  pottery,  and  porcelain.  The  presence  of  calcium 
and  magnesium  compounds  makes  the  clay  more  fusible,  because 
it  permits  the  formation  of  fusible  silicates  of  these  metals.  Bricks 
and  tiling  for  roofs  and  drains  are  made  of  common  clay  and,  when 
red,  owe  their  color  to  oxide  of  iron  Fe203.  The  firing  is  done  with 
fuel  gas  in  ovens  or  kilns  of  brickwork.  To  glaze  drain  pipes  and 
some  bricks,  salt  is  thrown  into  the  kiln.  The  water  vapor,  at  a 
red  heat,  hydrolyzes  the  salt,  hydrogen  chloride  is  set  free,  and  the 
sodium  hydroxide  gives  with  the  clay  a  fusible  sodium-aluminium 
silicate  which  fills  the  surface  pores.  Clay  for  fire  brick  (infusible) 
must  contain  silica,  but  110  lime. 

China  and  porcelain  are  made  from  pure  clay,  free  from  iron, 
to  which  a  little  of  the  more  fusible  felspar  is  added.  After  the 
first  firing,  the  articles  are  porous  (bisque),  and  must  be  covered 
with  a  glaze.  A  paste  of  finely  ground  felspar  and  silica,  some- 
times containing  lead  oxide,  is  spread  on  the  surface,  and  the 
articles  are  fired  again,  at  a  higher  temperature.  Colored  decora- 
tions are  added  by  means  of  suitable  materials,  mainly  oxides  of 
metals  which  give  colored  silicates.  The  third  firing  causes  these 
oxides  to  interact  and  fuse  with  the  glaze. 


562  COLLEGE    CHEMISTRY 

The  Schwerin  process  for  separating  ferric  oxide  FeiA  from 
clay,  so  that  white  porcelain  may  be  obtained,  is  now  used  on  a 
large  scale  and  affords  an  interesting  application  of  the  properties 
of  colloidal  suspensions  (p.  417).  When  the  impure  clay  is  sus- 
pended in  water,  the  particles  of  ferric  oxide  are  positively  charged 
and  those  of  the  clay  are  negative.  By  inserting  plates  connected 
with  the  dynamo  in  the  trough,  the  clay  particles  are  caused  to 
drift  towards  the  positive  plate  and  the  ferric  oxide  towards  the 
other,  so  that,  when  the  liquid  from  the  positive  end  is  allowed  to 
settle,  pure  clay  is  obtained. 

In  making  bricks,  in  some  cases,  advantage  is  taken  of  the  fact 
that  negative  colloids,  such  as  clay,  become  more  strongly  negative 
in  presence  of  a  trace  of  free  alkali.  Thus,  when  a  trace  of  sodium 
hydroxide  is  added  to  clay  slip,  the  particles  repel  one  another  more 
strongly,  the  cohesion  which  causes  the  plasticity  is  reduced,  and 
the  clay  can  be  poured  into  molds.  This  avoids  diluting  the  clay 
with  water,  which  would  only  have  to  be  driven  out  again,  with 
great  waste  of  heat,  in  the  firing. 

Cement.  —  Cement  is  made  by  heating  limestone  CaC03, 
clay  HAlSiO4,  and  sand  Si02,  or  a  natural  rock  containing  all  three 
in  the  right  proportions.  Such  a  rock,  made  into  cement  by 
volcanic  heat,  was  quarried  by  the  Romans  near  Naples  and  else- 
where, and  its  capacity  for  hardening  even  under  water  was  utilized 
by  them.  Blast-furnace  slag,  when  pulverized  and  heated  with 
limestone,  has  been  found  to  yield  an  excellent  quality  of  cement, 
and  a  valuable  use  has  thus  been  found  for  what  was  formerly  an 
annoying  encumbrarice.  The  mixture,  or  pulverized  natural  rock, 
is  moistened  and  fed  slowly  in  at  the  upper  end  of  an  inclined  (6°) 
revolving  cylinder  of  iron  (20  to  45  by  2  meters).  The  motion 
continually  turns  over  the  thin  layer,  and  exposes  every  particle 
to  the  heat  of  the  air-blast,  charged  with  pulverized  coal,  burning 
in  the  interior.  The  product  slides  out  in  a  continuous  stream  at 
the  lower  end,  and  is  pulverized  by  steel  balls  in  a  ball  mill. 

Cement  is  held  to  be  a  mixture  of  calcium  silicate  and  calcium 
aluminate.  The  former  is  simply  a  filler.  The  latter  is  hydrolyzed 
by  the  water: 

Ca3(AlO3)2  +  6H2O  -» 3Ca(OH)2  +  2H3A1O3. 


ALUMINIUM 


563 


The  calcium  hydroxide  slowly  crystallizes,  connecting  the  particles 
of  the  calcium  silicate.  The  aluminium  hydroxide  fills  the  inter- 
stices and  renders  the  whole  compact  and  impervious. 

Ultramarine.  —  Formerly,  pulverized  lapis  lazuli,  a  rare 
mineral  of  beautiful  blue  color,  was  used  by  artists  as  a  pigment. 
Gmelin  (1828)  found  a  way  of  making  it  artificially.  A  mixture  of 
kaolin,  sodium  carbonate,  sulphur,  and  charcoal  is  heated  until  a 
green  mass  is  obtained.  This  mass  is  then  pulverized  and  heated 
with  more  sulphur.  The  product  is  used  as  laundry  blueing,  in 
making  blue-tinted  paper,  and  with  oil  in  making  paint.  It  is  also 
added  in  small  amount  to  correct  the  natural  yellow  tint  of  linen, 
starch,  sugar  (p.  405),  and  paper-stock.  By  varying  the  mode  of 
heating,  without  altering  the  composition,  various  colors  from 
green  to  reddish  violet  can  be  obtained.  No  pure  colored  sub- 
stance can  be  extracted  from  it.  The  variety  of  colors  is  due  to 
different  degrees  of  colloidal  dispersion  of  some  substance  sus- 
pended in  the  solid,  just  as  gold,  which  is  pale  yellow  in  mass,  gives 
colloidal  suspensions  (p.  416)  of  different  colors  (red,  purple,  or 
blue)  according  to  the  fineness  of  the  particles  (cf.  p.  494). 

Dyeing.  —  The  problem  of  the  dyer  is  to  confer  the  desired 
color  upon  a  fabric  made,  usually,  of  cotton,  linen,  wool,  or  silk, 
and  to  do  this  in  stich  a  way  A  B 

that  the  dye  is  fast  to  (i.e.,  is 
not  removed  or  destroyed  by) 
rubbing  and  light,  and  often, 
also,  to  washing  with  soap.  To 
understand  the  means  by  which 
this  is  achieved,  it  must  be  noted 
that  cotton  and  linen  consist  of 
smooth  hollow  fibers  (Fig.  130A) 
of  the  composition  of  cellulose 
(C6Hi0O5)a;.  Wool  is  made  of  hollow  fibers,  with  a  scaley  surface 
(B)  and  silk  of  solid  filaments,  but  these  are  composed  of  proteins 
(p.  422).  Now,  the  proteins  are  much  more  active  chemically 
than  is  cellulose,  and  also,  as  colloidal  materials,  seem  to  have  a 
much  greater  tendency  to  adsorb  other  substances  than  has  cellu- 
lose. Hence,  accidental  stains  on  wool  or  silk  are  much  less  often 


FIG.  130. 


564  COLLEGE    CHEMISTRY 

removable  than  are  those  on  cotton,  and  when  samples  of  the  three 
materials  are  dipped  in  a  solution  of  a  dye,  the  first  two  are  per- 
manently dyed,  while  from  the  last  most  dyes  can  be  completely 
washed  out  with  water. 

Three  modes  of  dyeing  may  be  mentioned: 

1.  Insoluble  Dyes.     If  the  colored  body  can  be  produced  by 
precipitation,  after  the  solution  has  filled  the  capillary  and  wall 
of  every  fiber  of  the  goods,  then,  if  the  dye  is  sufficiently  insoluble, 
it  is  mechanically  imprisoned  in  every  fiber  and  cannot  be  washed 
out.     This  plan  may  be  applied  to  any  kind  of  goods.     For 
example,  if  cotton,  silk,  or  wool  is  first  boiled  in  a  solution  of  lead 
acetate,  and  is  then  soaked  in  a  boiling  solution  of  potassium 
chromate,  it  is  dyed  a  brilliant,  permanent  yellow.     Lead  chro- 
mate  is  the  colored  body: 

Pb(C02CH3)2  +  K2Cr04^  2K(C02CH3)  +  PbCr04| . 

The  part  precipitated  on  the  outside  of  the  goods  can  be,  and  is, 
at  once  washed  off  by  rubbing  in  water,  but  the  particles  inside  the 
fibers  can  come  out  only  by  being  dissolved,  and  they  are  insoluble 
in  water.  Indigo  Ci6Hi0N202,  which  is  used  in  larger  amounts  than 
any  other  dye,  belongs  to  this  class.  Obtained  in  early  times  from 
several  plants  in  Europe  and  Egypt,  where  it  was  known  as  woad, 
and  more  recently  imported  from  India,  where  the  cultivation  of 
the  indigo  plant  was  as  important  an  industry  as  is  the  growing 
of  cotton  in  the  Southern  States,  it  is  now  almost  all  made 
artificially.  Synthetic  indigo  is  manufactured,  with  naphthalene 
CioH8  (p.  411),  obtained  from  gas  tar  and  the  tar  from  by-product 
coke  ovens,  as  the  initial  substance.  The  cloth  is  saturated  with 
an  alkaline  solution  of  indigo  white  Ci6Hi2N2O2,  a  soluble,  slightly 
acid  substance,  and  the  oxygen  of  the  air  subsequently  oxidizes 
this  and  deposits  the  insoluble  indigo  blue  within  the  fibers: 
2C16H12N202  +  O2  -*  2C16H10N202 1  +  2H20. 

Indanthrene  blue  is  applied  in  the  same  way  as  indigo,  and  is  even 
less  affected  by  light. 

2.  Mordant  or  Adjective  Dyes.     Since  cotton  is  inactive  chemi- 
cally and,  although  a  colloid,  has  but  a  slight  tendency  to  adsorb 
dyes,  it  is  usually  necessary  first  to  introduce  into  the  fibers  of 
cotton  some  colloidal  substance  with  greater  adsorptive  powers. 
Substances  of  this  kind  are  tannic  acid  for  basic  dyes,  and  gelati- 


ALUMINIUM  ObO 

nous  colloidal  hydroxides,  such  as  those  of  aluminium,  tin,  iron  and 
chromium,  for  non-basic  (including  acid)  dyes.  They  are  called 
mordants  (Lat.  mordere,  to  bite).  Thus,  if  in  three  jars  we  place 
very  dilute  solutions  of  aluminium  sulphate,  ferric  chloride  FeCla 
and  chromous  acetate  Cr(C02CH3)2,  then  add  a  few  drops  of  a 
solution  of  a  dye  to  each,  and  finally  introduce  a  little  of  a  base 
(like  sodium  hydroxide)  to  precipitate  the  hydroxide  of  the  metal, 
this  hydroxide  will  adsorb  the  dye  and  carry  it  into  the  precipitate. 
Such  a  precipitate  of  mordant  and  dye  is  called  a  lake.  With  the 
same  dye,  the  three  lakes  have  different  colors.  Thus,  in  the  above- 
mentioned  experiment,  if  alizarin  (madder)  is  used  as  the  dye,  the 
colors  are  red  (Turkey  red),  violet,  and  maroon,  respectively. 
This,  of  course,  is  due  to  the  different  degrees  of  dispersion  in  the 
three  colloidal  materials.  If  aluminium  hydroxide  is  to  be  used, 
by  first  saturating  the  cloth  with  hot  aluminium  acetate  solution 
(p.  560),  or  by  using  first  aluminium  sulphate  and  then  ammonium 
hydroxide,  the  aluminium  hydroxide  is  precipitated  within  the 
fibers  of  the  goods.  When  the  material  is  then  dyed,  the  coloring 
matter  is  adsorbed  by  the  mordant,  with  which  it  forms  an  in- 
soluble lake,  within  the  fibers.  Basic  dyes,  like  Malachite  green 
and  Methylene  blue,  behave  similarly  with  tannic  acid,  or  an 
insoluble  salt  of  tannic  acid,  as  mordant.  It  will  be  seen  that,  so 
far  as  the  fabric  is  concerned,  this  process,  like  the  first,  is  a  mechan- 
ical one,  and  is  independent  of  the  chemical  nature  of  the  goods. 

3.  Direct  or  Substantive  Dyes.  Most  organic  dyes  are  direct 
dyes  on  silk  or  wool,  and  require  no  mordant  with  these  materials. 
The  actions  seem  to  be  sometimes  chemical,  but  more  often  cases 
of  adsorption  by  the  silk  or  wool  (both  colloids)  themselves.  A 
few  dyes  are  also  fast  on  cotton.  Congo-red  Na2C32H22N6S206  is 
fast  both  on  cotton  and  wool,  but  is  no  longer  much  used. 
Chrysophenin  is  now  one  of  the  commonest  dyes  of  this  class. 
These  dyes,  which  are  sodium  salts  of  complex  organic  acids,  are 
colloids  like  soap  (p.  417),  and  are  salted  out  within  the  fibers  of 
the  goods  by  adding  sodium  sulphate  to  coagulate  them  and  assist 
the  adsorption  by  the  cotton.  Once  adsorbed  in  this  way,  unlike 
soap,  they  cannot  be  washed  out. 

Analytical  Reactions  of  Aluminium  Compounds.  —  The 

alkalies,  and  alkaline  solutions  like  that  of  ammonium  sulphide, 


566  COLLEGE    CHEMISTRY 

precipitate  the  white  hydroxide.  The  product  is  soluble  in  excess 
of  the  active  alkalies.  Soluble  carbonates  also  throw  down  the 
hydroxide.  Aluminium  compounds,  when  heated  strongly  in  the 
flame  with  cobalt  salts,  give  a  blue  aluminate  of  cobalt  Co(A102)2. 

Exercises. —  1.   What  are  the  differences  between  zinc  and 
aluminium,  and  their  corresponding  compounds? 

2.  Construct  equations  showing,  (a)  the  hydrolysis  of  aluminium 
sulphate  (p.  558),  (6)  the  interaction  of  aluminium  sulphate  and 
cobalt  nitrate  in  the  Bunsen  flame. 

3.  Formulate  the  ionization  of  aluminium  hydroxide  (pp.  557, 
531). 

4.  Why  does  zinc  hydroxide,  in  spite  of  its  feebleness  as  a  base, 
dissolve  in  ammonium  hydroxide,  while  aluminium  hydroxide  does 
not? 


CHAPTER  XLI 
GERMANIUM,   TIN,   LEAD 

THE  metallic  elements  of  the  fifth  column  of  the  periodic  table 
are  germanium  (Ge,  at.  wt.  72.5),  tin  (Sn,  at.  wt.  119),  and  lead 
(Pb,  at.  wt.  207.2).  These  are  on  the  right  side,  while  titanium 
(Ti,  at.  wt.  48.1),  zirconium  (Zr,  at.  wt.  90.6),  cerium  (Ce,  at.  wt. 
140.25),  and  thorium  (Th,  at.  wt.  232.4)  occupy  the  left  side. 

The  Chemical  Relations  of  the  Family.  —  All  of  these  ele- 
ments show  a  maximum  valence  of  four.  Germanium,  tin,  and 
lead  are  also  bivalent.  In  this  respect  they  resemble  carbon  and 
differ  from  silicon,  which  is  more  closely  allied  to  the  elements  on 
the  left  side  of  the  column.  The  oxides  and  hydroxides  in  which 
these  three  elements  are  bivalent  become  more  basic,  and  the 
elements  themselves  more  metallic  in  chemical  relations,  with 
increase  in  atomic  weight.  In  this  they  resemble  the  potassium, 
calcium,  and  gallium  families.  Curiously  enough,  the  same  three 
hydroxides  are  also  acidic.  They  are  more  strongly  acidic  than 
is  zinc  hydroxide,  for  the  salts  they  form  by  interaction  with  bases 
are  less  hydrolyzed  than  are  the  zincates.  This  acidic  character 
likewise  increases  in  the  order  in  which  the  elements  are  named 
above. 

GERMANIUM 

Germanium  (p.  301)  forms  two  oxides  GeO  and  Ge02  correspond- 
ing to  those  of  carbon  and  of  tin.  Germanious  oxide  is  not  very 
definitely  basic  or  acidic,  and  the  sulphide  is  the  only  other  well- 
defined  compound  of  this  set.  Germanic  oxide  and  hydroxide  are 
acidic  entirely.  The  resemblance  to  carbon  is  shown  in  the  for- 
mation of  an  unstable  compound  with  hydrogen,  of  germanium 
chloroform  GeHCls  and  of  a  volatile  chloride  GeCU  (b.-p.  87°). 

TIN 

The  Chemical  Relations  of  the  Element.  —  Tin  is  both  biva- 
lent and  quadrivalent.  Each  of  the  oxides  and  hydroxides  SnO 

567 


568  COLLEGE    CHEMISTRY 

and  Sn(OH)2,  Sn02and  SnO(OH)2  (or  Sn(OH)4),  is  both  basic  and 
acidic,  so  that  there  are  really  four  series  of  compounds.  Still, 
stannous  hydroxide  is  mainly  a  base,  of  a  feeble  sort,  while  stannic 
hydroxide  is  mainly  an  acid.  Thus  we  have  stannous  chloride, 
sulphate,  and  nitrate,  which  are  stable,  although  they  are  all  more 
or  less  hydrolyzed  by  water,  and  sodium  stannite  Na2.Sn02  which 
is  unstable.  On  the  other  hand,  stannic  nitrate,  sulphate,  and 
chloride  are  completely  hydrolyzed  by  water,  while  sodium  stan- 
nate  Na^SnOa  is  comparatively  stable.  The  dioxide  Sn02  is  an 
infusible  solid,  resembling  silicon  dioxide.  Tin  has  a  tendency 
to  give  complex  acids  and  salts,  like  H2SnCl6,  (NH^.SnCle,  but 
these  are  ionized  also  to  a  small  extent  after  the  manner  of  double 
salts,  giving  ions  of  Sn++++.  Tin  forms  no  salts  with  weak  acids, 
like  carbonic  acid. 

Occurrence  and  Extraction.  —  Tin  has  long  been  in  use, 
specimens  of  it  being  found  in  Egyptian  tombs.  The  chief  ore  of 
tin  is  tinstone,  or  cassiterite  Sn02,  which  forms  square-prismatic 
crystals  whose  dark  color  is  due  to  the  presence  of  iron  compounds. 
The  ore  is  roughly  pulverized  and  washed,  to  remove  granite  or 
slate  with  which  it  is  mixed,  and  is  then  roasted,  to  oxidize  the 
sulphides  of  iron  and  copper,  and  drive  off  the  arsenic  which  it 
contains.  After  renewed  washing  to  eliminate  sulphate  of  copper 
and  oxide  of  iron,  it  is  reduced  with  coal  in  a  reverberatory  furnace. 
The  tin  is  afterwards  remelted  at  a  gentle  heat,  and  the  pure  metal 
flows  away  from  compounds  of  iron  and  arsenic.  The  metal  is 
produced  mainly  in  Banca  and  other  parts  of  the  East  Indies,  in 
Bolivia,  and  in  Cornwall. 

Physical  and  Chemical  Properties.  —  Tin  is  a  silver-white, 
crystalline  metal  of  low  tenacity  but  great  malleability  (tin-foil). 
Its  specific  gravity  is  7.3,  and  its  melting-point  about  232°. 

Tin  is  dimorphous  (p.  266).  In  1851,  the  tin  pipes  of  an  organ 
were  found  to  have  turned  largely  into  a  gray  powder.  In  1868  a 
shipment  of  blocks  of  tin  stored  in  the  custom  house  in  Petrograd 
was  found  to  have  changed  in  the  same  way.  Objects  of  tin  in 
museums  frequently  show  spots  indicating  the  presence  of  the 
"tin  pest,"  as  it  was  called.  It  now  appears  that  white,  metallic 
tin  is  stable  only  above  18°,  and  that  below  this  temperature  it  is 


TIN  569 

unstable  and  is  liable  to  change  into  gray  tin.  This  transition 
point  is  similar  to  that  of  sulphur  at  96°  (p.  265).  By  immersing 
the  tin  in  a  solution  of  pink-salt  (see  below),  the  change  is  hastened. 
When  the  two  kinds  of  tin  are  used  as  the  poles  of  a  cell,  and  are 
surrounded  by  pink-salt  solution,  no  difference  in  potential  is 
observed  at  18°.  But  below  18°,  white  tin,  being  unstable,  is 
more  active  and  becomes  positive,  while  above  18°,  gray  tin  be- 
comes positive. 

Tin-plate  is  made  by  dipping  carefully  cleaned  sheets  of  mild 
steel  into  molten  tin.  Vessels  of  copper  are  also  coated,  internally, 
with  tin,  to  prevent  the  formation  of  the  basic  carbonate  (p.  503). 
For  this  purpose  they  are  cleaned  with  ammonium  chloride, 
sprinkled  with  rosin  (to  reduce  the  oxide),  and  heated  to  230°. 
Molten  tin  is  then  spread  on  the  surface  with  a  piece  of  tow. 
Alloys  of  tin,  such  as  bronze  (p.  503),  soft  solder  (50  per  cent  lead), 
pewter  (25  per  cent  lead),  and  britannia  metal  (10  per  cent  anti- 
mony and  some  copper),  are  much  used  in  the  arts.  On  account 
of  the  action  of  soft  water  containing  dissolved  oxygen  on  lead 
(see  p.  574),  tin  pipes  are  preferred  for  distributing  distilled  water 
and  for  beer  pumps. 

Much  tin  is  now  recovered  by  treating  old  "tin  cans"  and  scrap 
tin-plate  with  dry  chlorine.  The  dried  gas  converts  the  tin  into 
stannic  chloride  SnCU,  which  is  used  to  make  mordants,  but  hardly 
attacks  the  iron  (p.  160).  The  process  is  called  de tinning. 

Tin,  although  it  displaces  hydrogen  from  dilute  acids,  is  not 
tarnished  by  moist  air.  With  warm  hydrochloric  acid  it  gives 
stannous  chloride  SnCl2  and  hydrogen.  Hot,  concentrated  sul- 
phuric acid  forms  stannous  sulphate  SnSQi  and  sulphur  dioxide 
(cf.  p.  276).  Nitric  acid,  when  cold  and  dilute,  interacts  with  it, 
giving  stannous  nitrate  Sn(N03)2,  and  a  portion  of  the  nitric  acid 
is  reduced  to  ammonia  (cf.  p.  354).  With  concentrated  nitric 
acid,  stannic  nitrate  is  formed,  but  most  of  this  salt  is  hydrolyzed 
by  the  water  at  the  high  temperature  of  the  action  (cf.  p.  535) , 
and  metastannic  acid  (H2SnO3)5  (/3-stannic  acid)  remains.  The 
final  result  is  shown  by  the  equation  (simplified) : 

Sn  +  4HN03  -*  H2Sn03  +  4N02  +  H2O. 

Tin  also  displaces  hydrogen  from  caustic  alkalies,  giving  a  meta- 
stannate,  such  as  sodium  metastannate  Na2Sn03. 


570  COLLEGE    CHEMISTRY 

Chlorides  of  Tin.  —  Stannous  chloride  SnCl2,2H20  is  made  by 
the  interaction  of  tin  and  hydrochloric  acid.  When  the  crystals 
are  heated,  or  when  a  strong  aqueous  solution  is  diluted,  the  salt  is 
partially  hydrolyzed.  In  the  latter  case  the  basic  chloride 
Sn(OH)Cl  is  deposited.  By  presence  of  excess  of  hydrochloric 
acid,  the  hydrolysis  is  prevented.  The  solution  is  used  as  a  mor- 
dant (p.  565). 

Stannous  chloride  tends  to  pass  into  stannic  chloride  SnCU,  and 
is  therefore  an  active  reducing  agent.  Thus,  it  reduces  the 
chlorides  of  mercury  (p.  534)  and  of  the  noble  metals,  liberating 
the  free  metals.  The  action  is  of  the  form  Hg++  +  Sn++  —  >  Hg 
+  Sn  '  '  '  '  .  It  also  reduces  free  oxygen,  or,  what  is  the  same  thing, 
is  oxidized  by  the  air.  In  this  case,  stannic  chloride  is  formed  in 
the  acid  solution  and  the  liquid  remains  clear;  in  the  neutral 
solution  a  precipitate  of  the  basic  chloride  is  formed  as  well: 

6SnCl2  +  2H2O  +  O2  -»  4Sn(OH)Cl  +  2SnCl4. 

Powdered  tin,  if  placed  with  the  acid  solution,  will  undo  the  effects 
of  this  action  by  reducing  the  stannic  salt  to  the  stannous  condition. 
When  chlorine  acts  upon  tin,  or  upon  stannous  chloride  (either 
solid  or  dissolved),  stannic  chloride  SnCU  is  formed.  The  com- 
pound is  a  colorless  liquid  (b.-p.  114°)  which  fumes  very  strongly 
in  moist  air,  giving  hydrochloric  acid  and  stannic  acid.  It  is 
almost  completely  hydrolyzed  by  water.  The  stannic  acid  which 
is  formed  is  not  precipitated,  however,  but  remains  in  colloidal 
suspension: 

4H20  <±  4HC1  +  Sn(OH)4. 


The  chloride,  with  small  amounts  of  water,  gives  hydrates,  of 
which  SnCl4,5H2O,  "oxymuriate  of  tin,"  is  used  as  a  mordant. 
Double  (or  perhaps  complex)  salts,  such  as  ammonium-stannic 
chloride  or  "  pink-salt"  (NH^SnCle  (used  as  a  mordant  on  cot- 
ton), are  readily  formed. 

Stannic  bromide  SnBr4  (b.-p.  201°)  resembles  stannic  chloride. 

d-Stannic  Acid  and  its  Salts.  —  When  a  solution  of  stannic 
chloride  is  treated  with  ammonium  hydroxide,  a  white,  gelatinous 
precipitate  of  a-stannic  acid  is  formed: 

+  4NH4OH  -*  4NH4C1  +  F2SnO3  +  H20. 


TIN  571 

The  precipitate  loses  water  gradually  until  the  dioxide  remains, 
and  neither  Sn(OH)4  nor  SnO(OH)2  is  obtainable  as  a  definite 
compound.  When  stannic  oxide  is  fused  with  caustic  soda,  sodium 
metastannate,  or  a-stannate  Na2Sn03,3H20,  is  formed: 

Sn02  +  2NaOH  ->  Na*Sn03  +  H2O. 

This  compound  is  used  as  a  mordant  under  the  name  of  "pre- 
paring salt."  When  its  solution  is  acidified,  a-stannic  acid,  the 
actual  mordant,  is  formed  by  double  decomposition.  This  a- 
stannic  acid  interacts  readily  with  acids  and  alkalies,  and  the 
chloride  obtained  from  it  is  identical  with  stannic  chloride  de- 
scribed above. 

Flannelette  and  other  cotton  goods  are  rendered  non-inflam- 
mable by  saturation  first  with  sodium  a-stannate  solution  and  then, 
after  drying,  with  ammonium  sulphate.  The  acid  is  too  feeble  to 
form  an  ammonium  salt: 

Na^SnOa  +  (NH4)2SO4  ->  Na2S04  +  SnO(OH)2  +  2NH3. 

The  sodium  sulphate  is  washed  out  and  the  goods,  after  being 
dried,  contain  stannic  oxide.  The  latter  cannot  afterwards  be 
removed  by  washing,  and  the  material  is  permanently  fireproof. 
Silk  is  also  loaded  with  stannic  oxide,  the  amount  used  varying 
from  25  to  300  per  cent  or  more. 

The  a-stannates  of  the  metals,  aside  from  those  of  potassium  and 
sodium,  like  the  silicates  and  carbonates  which  they  much  resemble, 
are  all  insoluble  in  water,  and  may  be  made  by  double  decompo- 
sition. 

^-Stannic  Acid,  or  Metastannic  Acid.  —  The  product  of  the 
action  of  nitric  acid  upon  tin  (p.  569)  is  a  hydrated  stannic  oxide 
like  the  foregoing  substance,  but  is  not  identical  with  it.  It  is  not 
easily  acted  upon  by  alkalies.  By  boiling  it  with  caustic  soda, 
however,  and  then  extracting  with  pure  water,  a  soluble  sodium 
p-stannate  Na2Sn5On  is  obtained.  /3-stannic  acid  is  also  very 
slowly  attacked  by  acids,  and  the  chloride  secured  from  it  is  not 
identical  with  the  ordinary  chloride.  For  these  reasons  it  is  sup- 
posed to  be  a  hydrate  of  a  polymer  of  stannic  oxide  (Sn02)s,- 
zH20.  When  fused  with  caustic  soda,  it  gives  the  same  a-stannate 
as  does  the  dioxide  itself. 


572  COLLEGE    CHEMISTRY 

The  Oxides  of  Tin.  —  When  stannous  oxalate  is  heated  in 
absence  of  air,  stannous  oxide  SnO  remains:  SnC204—  >SnO+C02 
+  CO.  It  is  a  black  powder  which  burns  in  the  air,  giving  the 
dioxide.  The  corresponding  hydroxide  Sn20(OH)2  is  formed  by 
adding  sodium  carbonate  to  stannous  chloride  solution.  It  is  a 
white  powder,  easily  dehydrated,  and  interacts  with  alkalies  to 
give  soluble  stannites,  such  as  Na2Sn02.  With  acids,  the  hydrox- 
ide gives  stannous  salts. 

Stannic  oxide  SnO2  is  found  in  nature  (p.  568),  and  may  be  made 
in  pure  form  by  igniting  0-stannic  acid.  When  heated,  it  becomes 
yellow,  but  recovers  its  whiteness  when  cooled  (cf.  Zinc  oxide,  p. 
528).  Prepared  at  a  low  temperature,  it  interacts  easily  with 
acids,  but  after  strong  ignition,  is  affected  by  them  very  slowly. 

The  Sulphides  of  Tin.  —  Stannous  sulphide  SnS  is  obtained  as 
a  dark-brown  precipitate  when  hydrogen  sulphide  is  led  into  a 
solution  of  a  stannous  salt. 

Stannic  sulphide  SnS2  is  formed  likewise  by  precipitation,  and  is 
yellow  in  color.  Stannic  sulphide  loses  sulphur  when  strongly 
heated,  and  leaves  stannous  sulphide.  It  is  not  much  affected  by 
dilute  acids,  but  interacts  with  solutions  of  ammonium  sulphide 
(or  sodium  sulphide),  giving  a  soluble  complex  sulphide,  namely, 
ammonium  sulphostannate  i 

SnS2  +  (NH4)2S  ->  (NH4)2.SnS3. 


The  corresponding  sodium  sulphostannate  is  easily  crystallized  in 
the  form  Na2SnS3,2H20.  Stannous  sulphide  is  not  affected  by 
soluble  sulphides,  but  polysulphides,  such  as  yellow  ammonium 
sulphide,  give  with  it  the  above-mentioned  sulphostannates  : 

SnS  +  (NH4)2S2  ->  (NH4)2.SnS3. 

With  acids  the  sulphostannates  undergo  double  decomposition, 
but  the  free  acid  H2.SnS3  thus  produced  is  unstable  and  breaks  up, 
giving  off  hydrogen  sulphide,  and  depositing  stannic  sulphide. 

Analytical  Reactions  of  Salts  of  Tin.  —  The  two  ionic 
forms  of  tin,  Sn++  and  Sn  '  '  '  '  ,  are  both  colorless.  Their  behavior 
is  different.  They  give  a  brown  and  a  yellow  sulphide,  respec- 
tively, with  hydrogen  sulphide.  These  sulphides  interact  with 


LEAD  573 

yellow  ammonium  sulphide  (above).  The  reducing  power  of 
stannous-ion  Sn++  is  very  characteristic  (p.  570).  The  oxides  are 
reduced  by  charcoal  in  the  reducing  part  of  the  Bunsen  flame  and 
the  metal  is  liberated. 

LEAD 

The  Chemical  Relations  of  the  Element.  —  Lead  is  both 
bivalent  and  quadrivalent.  The  oxides  PbO  and  Pb02,  and  the 
corresponding  hydrated  oxides,  are  all  both  basic  and  acidic.  Lead 
monoxide  is  a  fairly  active  base,  comparable  with  cupric  oxide,  but 
lead  dioxide  is  a  feeble  one.  Both  are  feebly  acidic.  The  salts  of 
bivalent  lead,  like  Pb(NO3)2,  commonly  called  the  plumbic  salts, 
are  somewhat  hydrolyzed  by  water,  but  less  so  than  are  those  of 
tin.  The  tetrachloride  and  other  salts  of  quadrivalent  lead  are 
completely  hydrolyzed.  The  plumbites  Na2.Pb02  and  plumbates 
Na2.Pb03,  like  the  stannites  and  stannates,  are  hydrolyzed  to  a 
considerable  extent.  All  the  compounds  in  which  lead  is  quad- 
rivalent give  up  half  of  the  negative  radical  readily,  and  are  re- 
duced to  the  "  plumbic  "  condition.  The  metal  displaces  hydrogen 
with  difficulty,  and  is  easily  displaced  by  zinc.  Lead  compounds 
are  all  poisonous,  and  the  effects  of  repeated,  very  minute  doses 
are  cumulative,  —  resulting  in  "lead  colic."  For  this  reason,  the 
manufacture  of  white  lead  is  forbidden  by  law  in  France,  and  is 
subject  to  strict  regulation  in  other  countries. 

Occurrence  and  Metallurgy.  —  Commercial  lead  is  almost  all 
obtained  from  galena  PbS,  which  crystallizes  in  cubes,  and  is  found 
in  the  United  States  (one-third  of  the  world's  supply),  Spain,  and 
Mexico.  This  ore  often  contains  considerable  amounts  of  silver 
sulphide  Ag2S  (cf.  p.  513). 

The  sulphide  of  lead  is  first  roasted  until  a  sufficient  proportion 
of  it  has  been  converted  into  the  oxide  and  sulphate.  The  furnace- 
doors  are  then  closed,  and  the  temperature  raised  in  order  that 
these  products  may  interact  with  the  unchanged  part  of  the 
sulphide:  pbs  +  2pb0  _^  3pb  +  SQ2j 

PbS  +  PbS04  ->  2Pb  +  2S02. 

Another  plan  consists  in  heating  galenite  with  scrap  iron  or  iron 
ores  and  coal:  PbS  +  Fe  — >  Pb  +  FeS.  The  molten  ferrous 
sulphide  rises  to  the  top  as  a  matte. 


574  COLLEGE    CHEMISTRY 

Lead  is  refined  elect rolytically  by  the  Betts  process.  Heavy 
plates  of  the  crude  lead  form  the  anodes,  thin  sheets  of  pure  lead 
the  cathodes,  and  a  solution  of  lead  fluosilicate  PbSiF6  the  cell 
liquid.  The  operation  is  similar  to  that  for  refining  copper  (p. 
511).  Silver,  gold  and  bismuth  are  left  as  a  sludge. 

Physical  and  Chemical  Properties.  —  Metallic  lead  is  gray 
in  color,  very  soft,  and  of  small  tensile  strength.  Its  specific 
gravity  is  11.4,  and  its  melting-point  327.4°.  While  warm,  it  is 
formed  by  hydraulic  pressure  into  pipes  which  are  used  in  plumbing 
and  for  covering  electric  cables.  On  account  of  its  very  slow  inter- 
action with  most  substances,  sheet  lead  is  used  in  chemical  fac- 
tories, for  example,  to  line  sulphuric-acid  chambers.  An  alloy 
containing  0.5  per  cent  of  arsenic  is  used  in  making  small  shot  and 
shrapnel  bullets.  Type-metal  contains  20-25  per  cent  of  antimony 
and  expands  on  solidifying,  giving  a  perfect  reproduction  of  the 
mold.  In  both  cases  greater  hardness  is  secured  by  the  addition 
of  the  foreign  metal.  Solder  contains  50  per  cent  of  tin  and,  being 
a  solution,  melts  at  a  low  temperature. 

Lead  oxidizes  very  superficially  in  the  air.  The  suboxide  Pb2O 
is  supposed  to  be  first  formed.  The  final  covering  is  a  basic  car- 
bonate. Contact  with  hard  waters  confers  upon  lead  a  similar 
coating  composed  of  the  carbonate  and  the  sulphate.  These  de- 
posits, being  insoluble  and  strongly  adherent,  enclose  the  metal 
and  protect  the  water  from  contamination  with  lead  compounds. 
Pure  rain-water,  however,  since  it  has  no  hardness,  and  contains 
oxygen  in  solution,  gives  the  hydroxide  Pb(OH)2,  which  is  notice- 
ably soluble.  Hence  lead  pipes  can  safely  be  used  only  with 
somewhat  hard  water.  When  heated  in  the  air,  lead  gives  the 
monoxide  PbO  or  minium  Pb304,  the  latter  at  lower  temperatures. 

The  metal  displaces  hydrogen  from  hydrochloric  acid  slowly. 
It  is  hardly  affected  by  cold  concentrated  sulphuric  acid  (cf.  p.  284). 
Nitric  acid  attacks  it  readily,  giving  lead  nitrate  and  oxides  of 
nitrogen  (p.  354). 

Chlorides  and  Iodide.  —  Plumbic  chloride  PbCl2  is  precipi- 
tated when  a  soluble  chloride  is  added  to  a  solution  of  a  lead  salt. 
It  is  slightly  soluble  in  water  (1.5  :  100)  at  18°,  and  much  more  so 
at  100°. 


LEAD  575 

Lead  tetrachloride  PbCU  is  a  solid  at  — 15°,  and  loses  chlorine  at 
the  ordinary  temperature.  It  is  made  by  passing  chlorine  into 
plumbic  chloride  suspended  in  hydrochloric  acid.  The  solution 
contains  H2PbCl6.  Ammonium  chloride  is  added  and  ammonium 
chloroplumbate  (NH4)2PbCl6  crystallizes  out.  When  this  is 
thrown  into  cold,  concentrated  sulphuric  acid,  an  oil,  PbCLi, 
settles  to  the  bottom.  The  oil  fumes  in  the  air,  and  closely  re- 
sembles stannic  chloride  SnCU.  With  little  water,  it  slowly  de- 
posits PbCl2  and  gives  off  chlorine.  With  much  water  it  is  quickly 
hydrolyzed,  and  lead  dioxide  is  thrown  down: 

PbCU  +  2H20  -»  Pb02  +  4HC1. 

The  yellow  lead  iodide  PbI2  is  formed  by  precipitation.  It  crys- 
tallizes in  yellow  scales  from  solution  in  hot  water. 

Oxides  and  Hydroxides.  —  There  are  five  different  oxides  of 
lead,  Pb20,  PbO,  Pb304,  Pb203,  and  Pb02.  The  suboxide  Pb2O  is  a 
dark-gray  powder,  formed  by  gently  heating  the  oxalate.  Plumbic 
oxide,  or  lead  monoxide  PbO,  is  made  by  cupellation  (p.  513)  of 
lead,  and  the  solidified,  crystalline  mass  of  yellowish-red  color  is 
sold  as  litharge.  All  the  other  oxides  yield  this  one  when  they 
are  heated  above  600°  in  the  air.  Plumbic  oxide  takes  up  carbon 
dioxide  from  the  air,  and  therefore  usually  contains  a  basic  car- 
bonate. The  oxide  is  used  in  making  glass  and  enamels  and  for 
preparing  salts  of  lead.  Mixed  with  glycerine,  it  gives  a  cement 
for  glass  or  stone. 

Plumbic  hydroxide  Pb(OH)2  is  formed  by  precipitation.  It 
gives  up  water  in  stages,  the  successive  products  being  Pb(OH)2, 
Pb2O(OH)2,  Pb302(OH)2.  These  substances  are  equivalent  in 
composition  to  PbO,H2O,  2PbO,H20,  and  3PbO,H2O  respectively. 
The  hydroxide  is  observably  soluble  in  water,  and  gives  a  solution 
with  a  faintly  alkaline  reaction.  With  acids  it  forms  salts  of  lead. 
It  interacts  also  with  potassium  and  sodium  hydroxides  to  form 
the  soluble  plumbites,  like  sodium  plumbite  Na2.Pb02. 

Minium,  or  red  lead,  Pb304,  gives  off  oxygen  when  heated: 
2Pb304  <=±  6PbO  +  02. 

On  account  of  unequal  heating  during  manufacture,  commercial 
red  lead  is  never  fully  oxidized,  and  always  contains  litharge. 
Conversely,  commercial  litharge  usually  contains  a  little  minium. 


576  COLLEGE    CHEMISTRY 

Minium,  when  heated  with  warm,  dilute  nitric  acid,  is  decom- 
posed, and  leaves  lead  dioxide  as  an  insoluble  powder.  Two- 
thirds  of  the  lead  is  basic  and  one-third  is  acidic.  Minium  is 
therefore  lead  orthoplumbate  (see  below) : 

Pb2.Pb04  +  4HN03  <=±  2Pb(N03)2  +  H4Pb04. 

The  double  decomposition  as  a  salt  that  it  thus  undergoes  is  fol- 
lowed by  dehydration  of  the  plumbic  acid,  which  is  unstable 
(HtPbC^  — » Pb02  +  2H20),  and  the  dioxide  remains.  Red  lead 
is  used  in  glass-making,  and,  when  mixed  with  oil,  gives  a  red  paint. 
Lead  dioxide  PbO2  may  be  obtained  as  described  above  in  the 
form  of  a  brown  powder.  It  is  usually  made  by  adding  bleaching 
powder  to  an  alkaline  solution  of  plumbic  hydroxide: 

Na2.PbO2  +  Ca(OCl)Cl  +  H20  -*  2NaOH  +  CaCl2  +  Pb02j. 

In  this  action  we  may  regard  the  free  lead  hydroxide,  formed  by 
hydrolysis  of  the  plumbite,  as  being  oxidized  by  the  bleaching 
powder.  Lead  dioxide  is  an  active  oxidizing  agent.  It  interacts 
with,  and  sets  fire  to,  a  stream  of  hydrogen  sulphide,  and  it  liber- 
ates chlorine  from  hydrochloric  acid.  With  acids  it  gives  no 
hydrogen  peroxide,  and  is  not  a  peroxide  (peroxidate)  in  the  re- 
stricted sense  of  the  term  (p.  223).  Lead  dioxide  interacts  with 
potassium  and  sodium  hydroxides,  giving  soluble  plumbates.  The 
potassium  salt  K2Pb03,3H2O  is  analogous  to  the  metastannate 
K2SnOs,3H2O  (p.  571).  A  mixture  of  calcium  carbonate  and  lead 
monoxide  absorbs  oxygen  when  heated  in  a  stream  of  air,  and  the 
yellowish-red  calcium  orthoplumbate  is  formed: 

4CaC03  +  2PbO  +  02  <±  2Ca2Pb04  +  4C02. 

The  action  is  reversible,  and  is  at  the  basis  of  Kassner's  method  of 
manufacturing  oxygen  from  the  air. 

Other  Salts  of  Lead.  —  Lead  nitrate  Pb(N03)2  may  be  made 
by  treating  lead,  lead  monoxide,  or  lead  carbonate  with  nitric  acid. 
It  forms  white,  anhydrous  octahedra.  The  nitrate  and  acetate 
(see  below)  are  the  salts  of  lead  which,  because  of  their  solubility 
(see  Table),  are  most  commonly  used.  On  account  of  hydrolysis, 
the  solution  of  the  nitrate  is  acid  in  reaction. 

Lead  carbonate  PbCO3  is  found  in  nature.  It  may  be  formed  as 
a  precipitate  by  adding  sodium  bicarbonate  to  lead  nitrate  solution. 


LEAD  577 

With  normal  sodium  carbonate,  a  basic  carbonate  Pb3(OH)2(CO3)2 
is  deposited.  This  basic  salt  is  identical  with  white  lead,  which  on 
account  of  its  superior  opacity,  has  better  covering  power  than 
zinc-white  (p.  528)  or  permanent  white  (p.  496).  The  substance 
is  manufactured  in  various  ways,  all  of  which  involve  the  oxidation 
of  the  lead  by  the  air,  the  formation  of  a  basic  acetate  by  the  inter- 
action of  vinegar  or  acetic  acid  with  the  oxide,  and  the  subsequent 
decomposition  of  the  salt  by  carbon  dioxide.  The  best  quality  is 
obtained  by  the  Dutch  method.  In  this,  gratings  of  cast  lead 
(" buckles")  are  placed  above  a  shallow  layer  of  vinegar  in  small 
pots.  These  pots  are  buried  in  manure,  which  by  its  decomposition 
furnishes  the  carbon  dioxide  and  the  necessary  warmth.  The  grat- 
ings are  gradually  converted  into  a  white  mass  of  the  basic  car- 
bonate. The  vapor  of  acetic  acid  arising  from  the  vinegar  may 
be  regarded  as  a  catalytic  agent,  since  it  is  used  over  and  over 
again. 

Lead  acetate  Pb^HsC^SH^O  is  made  by  the~action  of  acetic 
acid  on  litharge.  It  is  easily  soluble  in  water  and,  from  the  sweet 
taste  of  the  solution,  is  named  sugar  of  lead  (used  in  medicine). 
The  basic  salt  Pb(OH)(C2H302)  is  formed  by  boiling  a  solution  of 
lead  acetate  with  excess  of  litharge.  Unlike  most  basic  salts,  this 
basic  salt  is  soluble  in  water,  and  its  solution  has  a  faintly  alkaline 
reaction. 

Lead  sulphate  PbSC>4  occurs  in  nature  as  anglesite.  Being  insol- 
uble in  water,  it  is  easily  obtained  by  precipitation. 

Natural  lead  sulphide  PbS  (galena)  forms  black,  cubic  crystals 
with  a  silvery  luster.  The  precipitated  salt  is  amorphous.  It  is 
more  easily  attacked  by  active  acids  than  is  mercuric  sulphide  (cf. 
p.  531). 

The  Storage  Battery.  —  In  the  ordinary  lead  accumulator  the 

plates  consist  of  leaden  gratings.  The  openings  are  filled  with 
finely  divided  lead  in  one  plate  and  with  lead  dioxide  in  the  other. 
These,  and  the  dilute  sulphuric  acid  in  the  cell,  are  the  active  sub- 
stances when  the  cell  is  charged.  When  the  battery  is  used,  the 
SO4=  ions  migrate  towards  the  plates  filled  with  the  lead  (Fig.  131), 
and  convert  this  into  a  mass  of  the  insoluble  lead  sulphate: 
SO4=  +  Pb  ->  PbS04  +  20.  These  plates  receive  the  negative 
charges.  Simultaneously,  the  H+  ions  move  towards  the  other 


578 


COLLEGE    CHEMISTRY 


plates  and  there  reduce  to  monoxide  the  lead  dioxide  with  which 
they  are  filled. 

Pb02  +  2H+  -*  H20  +  PbO  +  20 . 

These  plates  acquire  positive  charges  and,  by  interaction  of  the 
lead  monoxide  with  the  sulphuric  acid,  become  filled,  like  the 
negative  plates,  with  lead  sulphate.  During  the  discharge,  much 
sulphuric  acid  is  thus  removed  from  the  cell  fluid,  and  the  approach- 
ing exhaustion  of  the  cells  can  thus  be  ascertained  by  measuring 
the  specific  gravity  of  the  fluid.  The  E.M.F.  of  the  current  is  a 
little  over  2  volts. 


PbOQ  PbS04 


DISCHARGE 


<-S04= 


CHARGE 


804= 


FIG.  131. 


Fia.  132. 


The  charging  is  done  by  passing  a  current  through  the  cell,  in 
the  opposite  direction  to  the  one  which  it  yields  (Fig.  132).  The 
H+  ions  are  attracted  to  the  negative  plate  and  an  equivalent 
number  of  SO4-  ions  are  formed,  so  that  only  lead  remains: 

PbS04  +  2H+  +  20  ->  Pb  +  2H+  +  S04= 

Simultaneously,  the  S04=  is  attracted  by  the  positive  plate  and, 
with  the  lead  sulphate  there  present,  forms  lead  persulphate: 
S04=  +  PbS04  +  2©  ->  Pb(S04)2.  The  persulphate,  being  a 
salt  of  quadrivalent  lead,  is  at  once  hydrolyzed  and  the  filling  of 
this  plate  is  thus  changed  into  lead  dioxide:  Pb(S04)2  +  2H2O  — > 
Pb02  +  2H2SO4.  Both  plates  are  thus  brought  back  to  the  con- 
dition in  which  they  were  before  the  discharge. 


LEAD  579 

The  last  set  of  charges  consumes  energy,  while  the  first  set 
liberates  energy.  Both  may  be  stated  in  a  single  equation: 

charge  — •> 

2PbS04  +  2H2O  «=*  Pb  +  2H2S04  +  Pb02. 
<—  discharge 

In  the  Edison  cell,  when  charged,  one  plate  is  of  iron  and  the 
other  contains  nickelic  oxide  Ni203.  The  cell  liquid  is  a  solution 
of  potassium  hydroxide.  When  the  cell  operates,  the  nickelic 
oxide  is  reduced  to  Ni(OH)2  and  the  iron  is  oxidized  to  Fe(OH)2, 
an  action  which  delivers  energy: 

Fe  +  3H2O  +  Ni203  <=±  Fe(OH)2  +  2Ni(OH)2. 

When  the  cell  is  charged,  the  nickel  is  reoxidized  and  the  iron 
reduced. 

Paints.  —  A  paint  usually  contains  three  ingredients: 

1.  The  oil  hardens  to  a  tough  resin  on  exposure  to  the  air 
("dries")  and  adheres  firmly  to  the  surface  being  painted. 

2.  The  body  is  a  fine  powder  which  makes  the  paint  opaque. 
Since  the  powder  does  not  shrink,  it  also  "fills"  the  paint  and  pre- 
vents the  formation  of  minute  pores  which  otherwise  would  appear 
in  the  oil  after  drying.     White  lead  (p.  577)  is  a  common  material 
for  the  body,  but  zinc  oxide,  lithopone  (p.  497)  and  other  sub- 
stances are  used. 

3.  Except  in  the  case  of  white  paint,  a  pigment  is  added.     Vari- 
ous oxides,  such  as  minium,  colored  salts,  and  lakes  (p.  565)  are 
used  as  coloring  matters. 

The  oil  does  not  "dry"  by  evaporation  but  gives  a  resin  by 
oxidation.  Linseed  oil  and  hemp  oil  are  commonly  used.  They 
contain  glyceryl  esters  (p.  414)  of  unsaturated  acids,  such  as  that 
of  linoleic  acid  C3H5(C02Ci7H3i)3,  which  contains  four  units  of 
hydrogen  less  than  stearic  acid.  The  unsaturated  part  of  the 
molecule  takes  up  the  oxygen.  By  previously  boiling  the  oil  with 
manganese  dioxide  and  other  oxides,  it  is  rendered  more  active, 
and  "dries"  more  quickly. 

Plumbers  use  a  cement  made  of  minium  and  linseed  oil,  in  which 
the  former  oxidizes  the  latter,  without  access  of  air  being  necessary. 


580  COLLEGE    CHEMISTRY 

Analytical  Reactions  of  Lead  Compounds.  —  Hydrogen 
sulphide  precipitates  the  black  sulphide,  even  when  dilute  acids 
are  present.  Sulphuric  acid  throws  down  the  sulphate.  Potas- 
sium hydroxide  gives  the  white  hydroxide,  which  dissolves  in 
excess  to  form  the  plumbite.  Potassium  chromate  or  dichromate 
(q.v.)  gives  a  yellow  precipitate  of  lead  chromate  PbCr04,  which  is 
used  as  a  pigment  under  the  name  of  "chrome-yellow." 

TITANIUM,  ZIRCONIUM,  CERIUM,  THORIUM 

The  metals  on  the  left  side  of  the  fifth  column  of  the  periodic 
table  are  all  quadrivalent,  although  compounds  in  which  a  lower 
valence  appears  are  numerous  in  this  family.  The  first  two  are 
feebly  base-forming  as  well  as  feebly  acid-forming;  the  last  two 
are  base-forming  exclusively. 

Titanium  occurs  in  rutile  TiTi04.  Derived  from  it  are  a  number 
of  titanates  of  the  form  K2TiOs.  Zirconium  is  found  in  zircon,  the 
orthosilicate  of  zirconium  ZrSK^.  The  oxide  is  used  in  making  the 
incandescent  substance  in  some  forms  of  gas  lamps. 

Cerium  occurs  chiefly  in  cerite  [Ce,  La,  Nd,  Pr]  Si04,H20  (cf.  p. 
443).  The  particles  of  an  alloy  of  cerium  (70  per  cent)  and  iron 
(30  per  cent),  when  torn  off  by  a  file,  catch  fire  in  the  air.  This 
fact  is  utilized  in  making  gas-lighters  and  cigar-lighters.  Thorium 
is  found  in  thorite  ThSi04,  but  most  of  the  supply  comes  from 
monazite  sand.  The  nitrate  Th(NO3)4,6H20  is  used  in  making 
Welsbach  incandescent  mantles  (cf.  Flame,  p.  397).  The  com- 
pounds are  radioactive  (see  Radium). 

The  foundation  of  the  Welsbach  mantle  is  woven  of  ramie.  This 
is  saturated  with  a  solution  of  thorium  and  cerium  nitrates  in  the 
proportion  99 : 1,  and  is  then  molded  to  the  proper  shape  and  dried. 
By  heating  in  a  Bunsen  flame,  the  organic  matter  is  burned,  and 
the  nitrates  are  decomposed: 

Th(N03)4  ->  Th02  -f  4  N02  +  02- 

The  oxides  retain  the  form  of  the  fabric  and,  to  prevent  breakage 
in  handling,  the  structure  is  dipped  in  collodion  and  dried. 

Exercises.  —  1.  In  what  order  should  you  place  the  elements 
dealt  with  in  this  chapter,  beginning  with  the  least  metallic,  and 
ending  with  the  most  metallic  (p.  436)? 


TITANIUM,    ZIRCONIUM,    CERIUM,    THORIUM  581 

2.  Construct  equations  showing,  (a)  the  interaction  of  tin  and 
concentrated  sulphuric  acid,  (6)  of  water  and.stannous  chloride,  (c) 
of  oxygen  and  stannous  chloride  in  acid  solution,  (d)  the  decom- 
position of  lead  oxalate  (p.  575),  (e)  the  interaction  of  lead  mon- 
oxide and  acetic  acid,  (/)  and  of  lead  monoxide  and  lead  acetate. 

3.  To  which  class  of  ionic  actions  (pp.  259,  270,  504)  do  the 
reductions  by  stannous  chloride  and  by  tin  (p.  570)  belong? 

4.  What  interactions  probably  occur  when  lead  dioxide  liberates 
chlorine  from  hydrochloric  acid? 

5.  How  should  you  set  about  preparing,  (a)  lead  oxalate  (in- 
soluble), (6)  lead  chlorate  (soluble)? 

6.  Construct  equations  for  the  formation  of  white  lead  by  the 
Dutch  process,  showing,  (1)  the  formation  of  the  basic  acetate  by 
the  action  of  oxygen,  water,  and  acetic  acid  vapor,  and  (2)  the 
action  of  carbonic  acid  on  the  product. 


CHAPTER  XLII 
ARSENIC,   ANTIMONY,   BISMUTH 

.THIS  family  is  very  closely  related  to  the  elements  phosphorus 
and  nitrogen  which  precede  it  in  the  same  column  of  the  periodic 
table.  In  reading  this  chapter,  therefore,  constant  reference 
should  be  made  to  the  chemistry  of  the  corresponding  compounds 
of  phosphorus.  For  a  general  comparison  of  the  elements  arsenic 
(As,  at.  wt.  75),  antimony  (Sb,  at.  wt.  120.2)  and  bismuth  (Bi,  at. 
wt.  208)  with  each  other  and  with  the  two  already  disposed  of,  see 
p.  592.  It  is  sufficient  here  to  say  that  arsenic  is  mainly  an  acid- 
forming  element,  and  is  therefore  a  non-metal,  while  antimony  is 
both  acid-forming  and  base-forming,  and  bismuth  is  base-forming. 
Each  of  the  three  elements  gives  two  sets  of  compounds,  in  which 
it  is  trivalent,  and  quinquivalent,  respectively.  None  of  the 
elements  when  free  displaces  hydrogen  from  dilute  acids. 

ARSENIC  As 

The  Chemical  Relations  of  the  Element.  —  Arsenic  forms  a 
compound  with  hydrogen  AsH3.  It  gives  several  halogen  deriva- 
tives of  the  type  AsX3,  which  are  completely  hydrolyzed  by  water. 
Its  oxides  and  hydroxides  are  acidic. 

Sulphates,  nitrates,  carbonates,  and  other  salts  of  arsenic  are  not 
formed.  The  complex  sulphides  (p.  572)  are  important. 

Occurrence  and  Preparation.  —  Arsenic  is  found  free  in 
nature.  It  occurs  also  in  combination  with  many  metals,  par- 
ticularly in  arsenical  pyrites  (mispickel)  FeAsS.  Two  sulphides 
of  arsenic,  orpiment  As2S3  and  realgar  As2S2,  and  an  oxide  As203, 
are  less  common. 

The  element  is  obtained  either  from  the  native  material  or  by 
heating  arsenical  pyrites :  FeAsS  — •>  FeS  -f-  As.  During  the  roast- 
ing of  the  sulphur  ores  of  metals,  arsenic  trioxide  is  formed  by  the 

582 


ARSENIC  583 

oxidation  of  the  arsenic  so  frequently  present,  and  collects  as  a  dust 
in  the  flues.  The  supply  is  greatly  in  excess  of  the  demand. 

Physical  and  Chemical  Properties.  —  The  free  element  is 
steel-gray  in  color,  metallic  in  appearance,  and  crystalline  in  form. 
It  gives  off  vapor  at  180°,  and  above  600°  acquires  a  vapor  pressure 
of  760  mm.  The  density  of  the  vapor  measured  at  644°  gives 
308.4  as  the  weight  of  the  G.M.V.  (22.4  liters  at  0°  and  760  mm.). 
The  weight  of  arsenic  combining  with  one  chemical  unit  weight 
(35.46  g.)  of  chlorine  is  25  g.  Three  times  this  amount,  or  75  g., 
is  the  smallest  weight  found  in  the  G.M.V.  of  any  volatile  com- 
pound of  arsenic,  and  is  therefore  accepted  as  the  atomic  weight 
(p.  106).  Since  308.4  is  equal  approximately  to  4  X  75  (=  300), 
the  formula  of  the  vapor  of  the  simple  substance  at  644°  is  As4.  At 
1700°  the  formula  is  As2  (cf.  p.  117). 

The  free  element  burns  in  the  air,  producing  clouds  of  the  solid 
trioxide  As203.  It  unites  directly  with  the  halogens,  with  sulphur, 
and  with  many  of  the  metals.  When  boiled  with  nitric  acid, 
chlorine- water,  and  other  powerful  oxidizing  agents  (p.  157),  it  is 
oxidized  in  the  same  way  as  is  phosphorus,  and  yields  arsenic  acid 
H3As04. 

Arsine  AsH3.  —  This  substance  corresponds  in  composition  to 
ammonia  and  phosphine,  and  some  of  the  ways  in  which  it  may  be 
formed  are  analogous  to  those  used  in  the  case  of  these  substances. 
Thus,  when  arsenic  and  zinc  are  melted  together  in  the  proportions 
to  form  zinc  arsenide  Zn3As2,  and  the  product  is  treated  with  dilute 
hydrochloric  acid,  the  result  is  similar  to  the  action  of  water  or  dilute 
acids  upon  calcium  phosphide,  and  arsine  is  evolved  as  a  gas: 

Zn3As2  +  6HC1  ->  2AsH3  +  3ZnCl2. 

Arsine  (arsenuretted  hydrogen)  is  formed  also  by  the  action  of 
nascent  hydrogen  (cf.  p.  360)  upon  soluble  compounds  of  arsenic. 
When  a  solution  of  arsenious  chloride  AsCl3  or  arsenic  acid  is  added 
to  zinc  and  hydrochloric  acid  in  a  generating  flask,  arsine  is  formed: 

AsCl3  +  3H2  -» AsH3  +  3HC1. 

Pure  arsine  may  be  secured  by  leading  the  mixture  with  hydrogen 
through  a  U-tube  immersed  in  liquid  air.  The  arsine  (b.-p.  —55°) 
condenses  as  a  colorless  liquid  (m.-p.  —119°). 


584  COLLEGE    CHEMISTRY 

Arsine  burns  with  a  bluish  flame,  producing  water  and  clouds  of 
arsenic  trioxide:  2AsH3  +  302  — »  3H2O  +  As2O3.  The  combus- 
tion of  hydrogen  containing  arsine,  generated  as  just  described, 
gives  the  same  substances.  Since  arsine,  when  heated,  is  readily 
dissociated  into  its  constituents  (cf.  p.  268),  the  vapor  of  free 
arsenic  is  present  in  the  interior  of  the  hydrogen  flame.  This 
arsenic  may  be  condensed  in  the  form  of  a  metallic-looking, 
brownish  stain  by  interposition  of  a  cold  vessel  of  white  porcelain 
(cf.  Fig.  85,  p.  268).  Even  when  only  a  trace  of  the  compound 
of  arsenic  has  been  added  to  the  materials  in  the  generator,  the 
stain  which  is  produced  is  very  conspicuous.  This  behavior  thus 
furnishes  us  with  the  basis  of  an  exceedingly  delicate  test  — 
Marsh's  test  —  for  the  presence  of  arsenic  in  any  soluble  form  of 
combination.  The  compounds  of  antimony  alone  show  a  similar 
phenomenon  (see  Stibine). 

Arsine  is  exceedingly  poisonous,  the  breathing  of  small  amounts 
producing  fatal  effects.  It  differs  from  ammonia  more  markedly 
than  does  phosphine,  for  it  is  not  only  without  action  on  water 
or  acids,  but  does  not  unite  directly  even  with  the  halides  of 
hydrogen. 

Halides  of  Arsenic.  —  The  halides  include  a  liquid  trifluoride 
AsF3,  a  liquid  trichloride,  a  solid  tribromide  AsBr3,  and  a  solid 
tri-iodide  AsI3. 

The  trichloride  AsCl3,  which  is  prepared  by  passing  chlorine  gas 
into  a  vessel  containing  arsenic,  is  easily  formed  as  the  result  of  a 
vigorous  action.  It  is  a  colorless  liquid  (b.-p.  130°).  When  mixed 
with  water  it  is  at  once* converted  into  the  white,  almost  insoluble 
trioxide.  The  action  is  presumably  similar  to  that  of  water  upon 
the  corresponding  compound  of  phosphorus  (p.  372),  but  the 
arsenious  acid  for  the  most  part  loses  water  and  forms  the  insoluble 
anhydride : 

AsCl3  +  3H20  <=>  As(OH)3  +  3HC1, 
2As(OH)3  <=±  As203 1  +  3H20. 

This  action,  however,  differs  markedly  from  the  other  in  that  it  is 
reversible,  and  arsenic  trioxide  interacts  with  aqueous  hydrochloric 
acid,  giving  a  solution  of  arsenious  chloride.  When  this  solution 
is  boiled,  arsenious  chloride  escapes  along  with  the  vapor. 


ARSENIC  585 

Oxides  of  Arsenic.  —  Arsenic  trioxide  As2Os  is  produced  by 
burning  arsenic  in  the  air  and  during  the  roasting  of  arsenical  ores 
(p.  582),  and  is  known  as  "white  arsenic"  or  simply  "arsenic." 
It  is  purified  for  commercial  purposes  by  subliming  the  flue-dust  in 
cylindrical  pots.  The  pure  trioxide  is  deposited  in  a  glassy  form 
in  the  upper  part  of  the  vessel.  Its  vapor  density  shows  it  to  have 
the  formula  As^e. 

When  treated  with  water,  the  trioxide  goes  into  solution  to 
slight  extent  (1.2  :  100  at  2°),  forming  arsenious  acid,  by  reversal  of 
the  second  of  the  actions  given  above.  In  boiling  water  the  solu- 
bility is  greater  (11.5  :  100).  When  heated  in  a  tube  with  carbon, 
this  oxide  is  reduced,  and  the  free  element,  being  volatile,  is  de- 
posited upon  the  cold  part  of  the  tube  just  above  the  flame.  The 
trioxide  is  an  active  poison,  since  it  gradually  passes  into  solution, 
forming  arsenious  acid.  The  fatal  dose  is  0.06-0.18  g.  (1-3  grains), 
but  "arsenic  eaters"  become  tolerant  of  it  and  can  take  four  times 
as  much  without  evil  effects. 

The  pentoxide  As20s  is  a  white  crystalline  substance,  formed  by 
heating  arsenic  acid :  2H3AsO4,H2O  — »  As2O5  +  4H20.  When  raised 
to  a  higher  temperature,  it  loses  a  part  of  its  oxygen,  leaving  the 
trioxide.  In  consequence  of  this  instability,  it  cannot  be  formed 
by  direct  union  of  oxygen  with  the  trioxide,  after  the  manner  of 
phosphorus  pentoxide. 

Acids  of  Arsenic.  —  When  elementary  arsenic  or  arsenious 
oxide  is  treated  with  concentrated  nitric  acid,  or  with  chlorine  and 
water,  orthoarsenic  acid  H3As04  is  produced.  The  substance 
crystallizes  as  a  deliquescent  white  solid  2H3As04,H20.  Salts  of 
this  acid,  and  of  pyroarsenic  acid  H4As2O7  and  metarsenic  acid 
HAsO3,  corresponding  to  the  phosphoric  acids  (p.  368),  are  known. 
The  two  last  acids,  themselves,  however,  are  not  known  as  such. 
It  has  been  shown  by  Menzies  that,  when  the  hemihydrate  of 
orthoarsenic  acid  is  dried  at  100°,  the  only  acid  obtainable  has  the 
composition  H5As3Oio(=  5H20,3As205).  When  this  acid  is  heated 
more  strongly,  it  loses  water,  leaving  the  pentoxide  As20s.  With 
metaphosphoric  acid,  the  final  elimination  of  all  the  water  by  simple 
heating  is  impossible.  The  chocolate-brown  silver  orthoarsenate 
Ag3As04  and  the  white  MgNH4As04,  like  the  corresponding  phos- 
phates, are  insoluble  in  water. 


586  COLLEGE    CHEMISTRY 

Arsenious  acid  HaAsOa,  like  sulphurous  and  carbonic  acids,  loses 
water,  and  yields  the  anhydride  (arsenic  trioxide)  when  the  attempt 
is  made  to  obtain  it  from  the  aqueous  solution.  The  potassium  and 
sodium  arsenites,  K3AsOa  and  NasAsOs,  are  made  by  treating  arsenic 
trioxide  with  caustic  alkalies,  and  are  much  hydrolyzed  by  water. 
The  arsenites  of  the  heavy  metals  are  insoluble,  and  can  be  made  by 
precipitation.  Paris  green  (p.  508)  is  an  arsenite  of  copper.  In 
cases  of  poisoning  by  white  arsenic,  freshly  precipitated  ferric 
hydroxide  (or  the  same  compound  in  colloidal  suspension)  or  mag- 
nesium hydroxide  is  administered,  since  by  interaction  with  the 
arsenious  acid  they  form  insoluble  substances. 


Sulphides  of  Arsenic.  —  Arsenic  pentasulphide  AsjjSs  is  ob- 
tained as  a  yellow  powder  by  decomposition  of  the  sulpharsenates 
(see  below),  and  by  leading  hydrogen  sulphide  into  the  solution  of 
arsenic  acid  in  concentrated  hydrochloric  acid  which  contains 
AsCl5. 

Arsenious  sulphide  As2Ss  occurs  in  nature  as  orpiment,  and  was 
formerly  used  as  a  yellow  pigment  (auripigmentum)  .  The  word 
arsenic  is  derived  from  the  Greek  name  for  this  mineral  (dpo-eviKm/). 
It  is  obtained  as  a  citron-yellow  precipitate  when  hydrogen  sul- 
phide is  led  into  an  aqueous  solution  of  arsenious  chloride. 

When  hydrogen  sulphide  is  led  into  an  aqueous  solution  of 
arsenious  acid,  the  sulphide  is  formed,  but  remains  in  colloidal 
suspension.  It  is  a  negatively  charged  colloid  (p.  417),  a  small 
amount  of  H+  ion  in  the  liquid  rendering  the  whole  electrically 
neutral.  It  is  coagulated  by  adding  solutions  of  salts,  lower  con- 
centrations being  sufficient  the  higher  the  valence  of  the  positive 
ion  of  the  salt  (0.05  Molar  KC1,  0.0007  M  BaCl2,  0.00009  M  A1C13). 

Realgar  As2S2  is  a  natural  sulphide  of  orange-red  color,  and  is  also 
manufactured  by  subliming  a  mixture  of  arsenical  pyrites  and 
pyrite: 

2FeAsS  +  2FeS2  -»  4FeS  +  As2S2  1  . 

It  burns  in  oxygen,  forming  arsenious  oxide  and  sulphur  dioxide, 
and  is  mixed  with  potassium  nitrate  and  sulphur  to  make  "  Bengal 
lights." 

Sulpharsenites  and  Sulpharsenates.  —  The  sulphides  of 
arsenic  interact  with  solutions  of  alkali  sulphides  after  the  manner 


ANTIMONY  587 

of  the  sulphides  of  tin  (p.  572),  giving  soluble,  complex  sulphides. 
Arsenious  sulphide  with  colorless  ammonium  sulphide  gives  ammo- 
nium sulpharsenite,  and  with  the  yellow  sulphide  gives  ammonium 
sulpharsenate  : 

3(NH4)2S  +  As2S3  ->  2(NH4)3.AsS3, 
3(NH4)2S  +  As2S3  +  2S  ->  2(NH4)3.AsS4. 

Proustite  (p.  512)  is  a  natural  sulpharsenite  of  silver. 

These  salts  are  decomposed  by  acids,  and  give  the  feebly  ionized 
sulpharsenious  or  sulpharsenic  acid: 

(NH^a.AsSs  +  3HC1  -*  3NH4C1  +  H3AsS3  ->  3H2S  t  +  As&l, 
(NH4)3.AsS4  +  3HC1  -+.  3NH4C1  +  H3AsS4  ->  3H2S  f  +  As2S5i  . 

These  sulpho-acids,  however,  at  once  break  up,  giving  hydrogen 
sulphide  as  a  gas,  and  the  sulphides  of  arsenic  as  yellow  precipitates. 

ANTIMONY  Sb 

The  Chemical  Relations  of  the  Element.  —  Antimony 
resembles  arsenic  in  forming  a  hydride  SbH3  and  halides  of  the 
forms  SbX3  and  SbXs.  The  latter  are  partially  hydrolyzed  by 
water  with  ease,  but  complete  hydrolysis  is  difficult  to  accomplish 
with  cold  water.  The  oxide  Sb2O3  is  basic  and  also  feebly  acidic 
(amphoteric),  and  the  oxide  Sb2(>5  is  acidic.  The  compositions  of 
the  compounds  are  similar  to  those  of  the  compounds  of  arsenic, 
but  there  are  in  addition  salts,  such  as  Sb2(S04)3,  derived  from  the 
oxide  Sb20s.  The  element  gives  complex  sulphides. 

Occurrence  and  Preparation.  —  Antimony  occurs  free  in 
nature.  The  black  trisulphide  Sb2S3,  stibnite,  is  found  in  Hungary 
and  Japan,  and  forms  shining,  prismatic  crystals.  Stibnite  is 
roasted  in  the  air  in  order  to  remove  the  sulphur,  and  the  white 
oxide  which  remains  is  mixed  with  carbon  and  reduced  by  strong 
heat: 

Sb2S3  +  502  -»  Sb2O4  +  3S02, 
Sb204  +  4C-»2Sb 


Properties.  —  Antimony  is  a  white,  crystalline  metal,  melting 
at  630°  (b.-p.  1300°)  .  It  is  brittle,  and  easily  powdered.  Its  vapor 
at  1640°  has  the  formula  Sb2,  while  at  lower  temperatures  Sb4  is 


588  COLLEGE    CHEMISTRY 

present.  It  is  used  in  making  alloys  such  as  type-metal,  stereo- 
type-metal, and  britannia  metal  (q.v.).  The  alloys  of  antimony 
expand  during  solidification,  and  therefore  give  exceptionally  sharp 
castings. 

Babbitt's  Metal  (Sb  3,  Zn  69,  As  4,  Pb  5,  Sn  19),  and  other  anti- 
friction alloys  used  in  lining  bearings,  contain  antimony  along  with 
zinc,  copper,  and  other  metals.  Molten  mixtures  of  metals 
(alloys),  when  solidifying,  do  not  always  form  a  homogeneous, 
solid  mass.  In  an  anti-friction  alloy,  what  is  wanted  is  a  mass,  in 
general  soft,  but  containing  hard  particles.  The  latter  bear  most 
of  the  pressure,  yet,  as  the  alloy  wears,  they  are  pressed  into  the 
softer  matrix  so  that  a  smooth  surface  is  always  presented.  An 
alloy  which  has  the  opposite  composition,  that  is,  which  gives  a 
hard  mass  containing  softer  particles,  develops  heat  by  friction 
much  more  rapidly. 

The  element  unites  directly  with  the  halogens.  It  does  not  rust, 
but  when  heated  it  burns  in  the  air,  forming  the  trioxide  Sb203  or  a 
higher  oxide  Sb204.  When  heated  with  nitric  acid,  it  yields  the 
trioxide  and,  with  more  difficulty,  antimonic  acid  H3Sb04. 

Stibine  SbHs.  —  The  hydride  of  antimony  SbH3  is  formed  by 
the  action  of  zinc  and  hydrochloric  acid  on  any  soluble  compound 
of  antimony.  By  the  action  of  dilute,  cold  hydrochloric  acid  on 
an  alloy  of  antimony  and  magnesium  (1  :  2),  a  mixture  of  hydrogen 
and  stibine  containing  as  much  as  11.5  per  cent  (by  volume)  of  the 
latter  may  be  made.  It  is  separated  by  cooling  with  liquid  air 
(b.-p.  —17°,  m.-p.  —88°).  It  is  more  easily  dissociated  than  is 
arsine  (p.  584),  and  forms  a  deposit  of  antimony  when  a  porcelain 
vessel  is  held  in  the  flame. 

Antimony  Halides.  —  The  halides  include  the  trichloride;  the 
pentachloride  SbCl6,  a  liquid  (m.-p.  -6°,  b.-p.  140°);  the  tribro- 
mide  SbBr3,  tri-iodide  SbI3,  trifluoride  SbF3,  and  pentafluoride  SbF5. 

Antimony  trichloride  SbCls  is  made  by  direct  union  of  chlorine 
and  antimony.  It  forms  large,  soft  crystals  (m.-p.  73°,  b.-p.  223°), 
and  used  to  be  named  "  butter  of  antimony."  When  treated  with 
little  water,  it  forms  a  white,  opaque,  insoluble  basic  salt,  anti- 
mony oxy  chloride: 

SbCl3  +  H20  <=±  SbOCl  |  +  2HC1. 


ANTIMONY  589 

With  a  large  amount  of  water,  a  greater  proportion  of  the  chlorine 
is  removed,  and  Sb405Cl2  (=  2SbOCl,Sb2O3)  remains.  With 
boiling  water  the  oxide  is  finally  formed.  The  action  is  not  com- 
plete as  long  as  hydrochloric  acid  is  present.  It  may  therefore  be 
reversed,  so  that,  on  addition  of  hydrochloric  acid  to  the  mixture, 
a  clear  solution  of  the  trichloride  is  re-formed.  If  the  concentra- 
tion of  the  acid  is  once  more  reduced  by  dilution  with  water,  the 
oxychloride  is  again  precipitated. 

Oxides  of  Antimony.  —  The  trioxide  Sb203  (vapor  density 
gives  Sb^e)  is  obtained  by  oxidizing  antimony  with  nitric  acid,  or 
by  combustion  of  antimony  with  a  limited  supply  of  oxygen.  It  is 
a  white  substance,  insoluble  in  water.  It  is  in  the  main  a  basic 
oxide,  interacting  with  many  acids  to  form  salts  of  antimony. 
But  it  interacts  also  with  alkalies,  giving  soluble  antimonites. 
The  pentoxide  Sb205  is  a  yellow,  amorphous  substance,  obtained 
by  heating  antimonic  acid.  It  combines  only  with  bases  to  form 
salts,  and  is  therefore  an  acid-forming  oxide  exclusively.  •  The 
tetroxide  Sb204  is  formed  by  heating  antimony  or  the  trioxide  in 
excess  of  oxygen.  It  is  neither  acid-  nor  base-forming. 

Salts  of  Antimony.  —  The  nitrate  Sb(N03)3  and  the  sulphate 

Sb2(S04)3  are  made  by  the  interaction  of  the  trioxide  with  nitric  and 
sulphuric  acids.  They  are  hydrolyzed  by  water,  giving  basic  salts, 
such  as  (SbO)2S04  (=  Sb2O2S04),  which,  like  SbOCl,  are  derived 
from  the  hydroxide  SbO(OH).  When  the  trioxide  is  heated  with 
a  solution  of  potassium  bitartrate  KH^HOe,  a  basic  salt  K(SbO)- 
C4H406,^H20,  known  as  tartar-emetic,  is  formed.  This  is  a  white, 
crystalline  substance  which  is  soluble  in  water  and  is  used  in 
medicine.  The  univalent  group  SbO1  is  known  as  antimonyl,  and 
the  above  mentioned  basic  compounds  are  often  called  antimonyl 
sulphate,  etc. 

Antimonic  Acid.  —  By  vigorous  oxidation  of  antimony  with 
nitric  acid,  or  by  decomposing  the  pentachloride  with  water,  a 
white,  insoluble  substance  of  the  approximate  composition  H3SbO4 
is  obtained.  This  substance  interacts  with  caustic  potash  and 
passes  into  solution.  But  the  salts  which  have  been  made  are 
pyro-  and  metantimoniates.  Thus,  when  antimony  is  fused  with 


590  COLLEGE    CHEMISTRY 

niter,  potassium  metantimoniate  KSb03  is  formed.  When  dis- 
solved, this  salt  takes  up  water,  giving  a  solution  of  the  acid 
potassium  pyroantimoniate: 

2KSbO3  +  H20  ->  K2H2Sb2O7. 

If  this  is  added  to  a  strong  solution  of  a  salt  of  sodium,  an  acid 
sodium  pyroantimoniate  is  thrown  down,  Na2H2Sb2O7.  This  is 
almost  the  only  somewhat  insoluble  salt  of  sodium. 

Sulphides  of  Antimony.  —  The  trisulphide  Sb2S3  is  found  in 
nature  as  the  black,  crystalline  stibnite.  As  precipitated  from 
solutions  of  salts  of  antimony,  the  trisulphide  is  an  orange-red 
powder,  which,  however,  after  being  melted,  assumes  the  appearance 
of  stibnite : 

2SbCl3  +  3H2S  <=±  Sb2S3 1  +  6HC1. 

Antimony  trisulphide,  like  cadmium  sulphide  (p.  531),  cannot  be 
precipitated  in  presence  of  concentrated  hydrochloric  acid. 

The  pentasulphide  Sb2S5  is  obtained  by  the  decomposition  of 
the  sulphantimoniates  (see  below).  In  appearance  it  resembles  the 
trisulphide  and,  when  heated,  decomposes  into  this  substance  and 
free  sulphur. 

The  sulphides  of  antimony  behave  towards  solutions  of  the 
alkali  sulphides  as  do  the  sulphides  of  arsenic  (p.  587).  The  tri- 
sulphide dissolves  in  colorless  ammonium  sulphide  with  difficulty, 
forming  an  unstable,  soluble  ammonium  sulphantimonite 

Sb2S3  +  3(NH4)2S  -» 2(NH4)3SbS3. 

With  the  pentasulphide  or  with  yellow  ammonium  sulphide  the 
soluble  ammonium  sulphantimoniate  is  readily  formed: 

Sb2S6  +  3(NH4)2S  -*  2(NH4)3.SbS4, 
Sb2S3  +  3(NH4)2S  +  2S  -»  2(NH4)3.SbS4. 

The  most  familiar  substance  of  this  class  is  Schlippe's  salt  Na3SbS4, 
9H20.  Pyrargyrite  (p.  512)  is  a  natural  sulphantimonite. 

When  acids  are  added  to  solutions  of  sulphantimoniates,  the 
sulphantimonic  acid  which  is  liberated  decomposes,  and  antimony 
pentasulphide  is  thrown  down  (see  under  Arsenic,  p.  587). 


BISMUTH  591 

BISMUTH 

The  Chemical  Relations  of  the  Element.  —  Bismuth  forms 
no  compound  with  hydrogen.  Its  compounds  with  the  halogens 
are  of  the  form  BiX3  and  are  hydrolyzed  by  water  giving  basic  salts. 
The  oxide  Bi203  is  basic,  and  the  oxide  Bi2O5  is  not  acidic.  Bis- 
muth gives  a  carbonate,  nitrate,  phosphate,  and  other  salts,  in 
which  it  acts  as  a  trivalent  element.  It  forms  no  soluble  complex 
sulphides. 

Occurrence  and  Properties.  —  This  element  is  found  free  in 
nature,  and  also  as  trioxide  Bi203  and  trisulphide  Bi2S3.  It  is  a 
shining,  brittle  metal  with  a  reddish  tinge  (m.-p.  271°).  Bismuth 
is  one  of  the  few  substances  (see  water)  which  expand  on  solidify- 
ing, the  crystals  being  lighter  than  the  liquid  at  271°.  It  is  di- 
morphous, with  a  transition  point  (p.  86)  at  75°.  Mixtures  of 
bismuth  with  other  metals  of  low  melting-point  fuse  at  lower 
temperatures  than  do  the  separate  metals.  This  is  a  corollary  of 
the  fact  that  a  solution  freezes  at  a  lower  temperature  than  does 
the  pure  solvent  (p.  134).  Thus,  Wood's  metal,  containing  bis- 
muth (m.-p.  271°)  4  parts,  lead  (m.-p.  327°)  2  parts,  tin  (m.-p. 
232°)  1  part,  and  cadmium  (m.-p.  321°)  1  part,  melts  at  60.5°, 
considerably  below  the  boiling-point  of  water.  Similar  alloys 
are  used  for  safety  plugs  in  steam-boilers  and  automatic 
sprinklers. 

Bismuth  does  not  tarnish,  but  when  heated  strongly  it  burns  to 
form  the  trioxide.  With  the  halogens  it  forms  a  fluoride  BiF3,  a 
bromide  BiBr3,  and  an  iodide  BiI3.  When  the  metal  is  treated 
with  oxygen  acids,  or  the  trioxide  with  any  acids,  salts  are  pro- 
duced. 

Compounds  of  Bismuth.  —  In  addition  to  the  basic  trioxide 
Bi203,  which  is  a  yellow  powder  obtained  by  direct  oxidation  of  the 
metal  or  by  ignition  of  the  nitrate,  three  other  oxides  are  known  — 
BiO,  Bi204,  and  Bi205.     None  of  these,  however,  is  either  acid- 
forming  or  base-forming. 

The  salts  of  bismuth,  when  dissolved  in  water,  give  insoluble 
basic  salts,  and  the  actions  are  reversible,  the  basic  salts  being 
redissolved  by  addition  of  an  excess  of  the  acid.  In  the  case  of  the 


592  COLLEGE    CHEMISTRY 

chloride  BiCl3,H2O  and  the  nitrate  Bi(NO3)3,5H2O,  the  actions 
taking  place  are: 

BiCU  +  2H20  <=*  Bi(OH)2Cl  +  2HC1, 
Bi(N03)3  +  2H20  <=*  Bi(OH)2NO3  +  2HNO3. 

The  former  of  these  products,  when  dried,  loses  a  molecule  of  water, 
giving  the  oxychloride  BiOCl.  The  oxynitrate  Bi(OH)2NO3  is 
much  used  in  medicine,  for  the  treatment  of  some  forms  of  in- 
digestion, under  the  name  of  "subnitrate  of  bismuth."  It  is  often 
contained  in  face  powders. 

The  brownish-black  trisulphide  Bi2S3  may  be  obtained  by  direct 
union  of  the  elements,  or  by  precipitation  with  hydrogen  sulphide. 
This  sulphide  is  not  affected  by  solutions  of  ammonium  sulphide  or 
of  potassium  sulphide.  It  differs,  therefore,  markedly  from  the 
sulphides  of  arsenic  and  antimony  in  its  behavior. 

THE  FAMILY  AS  A  WHOLE 

The  elements  themselves  change  progressively  in  physical 
properties  as  the  atomic  weight  increases.  Nitrogen  is  a  gas 
which  with  sufficient  cooling  yields  a  white  solid,  phosphorus  an 
almost  white  or  a  red  solid,  and  arsenic,  antimony,  and  bismuth 
are  metallic  in  appearance.  The  first  combines  directly  with  hy- 
drogen, the  next  three  give  hydrides  indirectly,  and  the  last  does 
not  unite  with  hydrogen  at  all.  The  hydride  of  nitrogen  combines 
with  water  to  form  a  base,  while  the  other  hydrides  show  no  such 
tendency.  Ammonia  unites  with  acids,  including  those  of  the 
halogens,  to  form  salts;  phosphine  with  the  hydrogen  halides  only; 
the  others  do  not  combine  with  acids  at  all.  As  regards  their 
metallic  properties,  in  the  chemical  sense,  nitrogen  and  phosphorus 
do  not  by  themselves  form  positive  ions,  and  furnish  us  therefore 
with  no  salts  whatever.  Arsenic  gives  a  trivalent  positive  ion, 
which  is  found  in  solutions  of  the  halides  only.  It  forms  no  normal 
sulphates,  nitrates,  or  other  salts.  Antimony  and  bismuth  both 
give  trivalent  positive  ions.  The  sulphates,  nitrates,  etc.,  of 
antimony,  however,  are  readily  decomposed  by  water  with  pre- 
cipitation of  the  hydroxide*.  The  salts  of  bismuth,  on  the  other 
hand,  do  not  readily  give  the  pure  hydroxide  with  water,  although 
they  are  easily  hydrolyzed  to  basic  salts. 


VANADIUM,    COLUMBIUM,    TANTALUM  593 

The  halogen  compounds  of  nitrogen  and  phosphorus  are  com- 
pletely hydrolyzed  by  water,  and  do  not  persist  when  any  water  is 
present,  even  when  excess  of  the  halogen  acid  is  used.  The  halogen 
compounds  of  arsenic  are  completely  hydrolyzed  by  cold  water,  but 
exist  in  solution  in  presence  of  excess  of  the  acids.  The  halogen 
compounds  of  antimony  and  bismuth  are  incompletely  hydrolyzed 
by  cold  water. 

Each  element  gives  a  trioxide  and  a  pentoxide.  With  nitrogen 
these  are  acid-forming,  being  the  anhydrides  of  nitrous  and  nitric 
acids.  With  phosphorus  the  trioxide  and  the  pentoxide  are  an- 
hydrides of  acids.  With  arsenic  the  trioxide  is  basic  towards  the 
halogen  acids,  and  is  the  first  example  of  a  basic  oxide  which  we 
encounter  in  this  group.  The  pentoxide,  however,  is  acid-forming. 
The  trioxide  of  antimony  is  mainly  base-forming,  although  it  is 
feebly  acid-forming  also.  The  pentoxide  is  acid-forming.  The 
trioxide  of  bismuth  is  base-forming  exclusively,  and  the  pentoxide 
has  no  derivatives. 

These  statements,  which  could  easily  be  expanded,  are  sufficient 
to  show  that  when  the  periodic  law  is  borne  in  mind  it  furnishes 
valuable  aid  in  systematizing  the  chemistry  of  a  group  like  this. 

Analytical  Reactions  of  Arsenic,  Antimony,  and  Bismuth. 

—  The  ions  which  are  most  frequently  encountered  are  AS+++, 
Sb+++,  Bi+++,  As04=-,  and  As03=-.  The  first  three,  with  hydro- 
gen sulphide,  give  colored  sulphides  which  are  not  affected  by 
dilute  acids.  The  sulphides  of  arsenic  and  antimony  are  separable 
from  the  sulphide  of  bismuth  by  solution  in  yellow  ammonium 
sulphide.  Marsh's  test  enables  us  to  recognize  the  presence  of 
traces  of  compounds  of  arsenic  and  antimony.  Oxygen  compounds 
of  arsenic,  when  heated  with  carbon,  give  a  volatilijpie^allic- 
looking  deposit  of  arsenic. 

VANADIUM,  COLUMBIUM,  ^TANTALUM 

Of  these  elements,  vanadium  is  less  uncommon  than  the  others. 
It  is  found  in  rather  complex  compounds.  When  these  are  J^eated 
with  soda  and  sodium  nitrate,  sodium  metavanadate  NaVOa  is 
formed,  and  can  be  extracted  with  lBH'tTpne  element  forms 
several  chlorides,  such  as  VC12,  VC13,  VTVvOfcl3,  and  five  oxides, 
V2O,  VO,  V203,  V02,  and  V2O5.  The  element  has  very  feeble  base- 


594 


COLLEGE    CHEMISTRY 


forming  properties,  and  gives  only  a  few  unstable  salts.     Ferro- 
vanadium,  an  alloy,  is  used  in  making  vanadium  steel. 

Columbium  (or  niobium),  first  discovered  and  named  by  Hat- 
chett  (1801),  and  tantalum  possess  feeble  base-forming  properties, 
their  chief  compounds  being  the  columbates  and  tantalates. 

Exercises.  —  1.  How  do  you  account  for  the  fact  that  the 
molecular  weight  of  arsenic  at  644°  is  not  exactly  300,  and  why  is 
308.4  -s-  4  not  accepted  as  the  atomic  weight? 

2.  Formulate  the  series  of  changes  involved  in  the  solution  of 
arsenic  trioxide  and  the  interaction  of  hydrochloric  acid  with  the 
arsenious  acid  so  formed  (cf.  p.  272). 

3.  What  is  the  full  significance  of  the  fact  that  arsenic  -penta- 
sulphide  may  be  precipitated  by  hydrogen  sulphide  from  a  solution 
of  arsenic  acid  in  hydrochloric  acid?     Make  the  equation. 

4.  To  what  classes  of  chemical  changes  do  the  interactions  of 
arsenious  sulphide  and  antimony  trisulphide  with  yellow  ammo- 
nium sulphide  belong? 

5.  Construct  equations  showing  the  interaction  of,  (a)  oxygen 
and  arsenical  pyrites,  (b)  chlorine-water  and  arsenic,  (c)  the  de- 
hydration of  orthoarsenic  acid,  (d)  potassium  hydroxide  and  arsenic 
trioxide,  (e)  concentrated  nitric  acid  and  antimony,  (/)  potassium 
bitartrate  and  antimony  trioxide,  (g)  acids  and  ammonium  ortho- 
sulphantimoniate. 

6.  How  should  you  set  about  making  Schlippe's  salt? 


CHAPTER  XLIII 
THE   CHROMIUM  FAMILY.    RADIUM 

THE  chromium  (Cr,  at.  wt.  52)  family  includes  molybdenum  (Mo, 
at.  wt.  96),  tungsten  (W,  at.  wt.  184),  and  uranium  (U,  at.  wt. 
238.2),  and  occupies  the  seventh  column  of  the  periodic  table  along 
with  the  sulphur  and  selenium  family. 

The  Chemical  Relations  of  the  Family.  —  The  features 
which  are  common  to  the  four  elements  are  also  those  which 
affiliate  them  most  closely  with  their  neighbors  on  the  right  side 
of  the  column.  They  yield  oxides  of  the  forms  CrOa,  Mo03, 
WO3,  and  U03,  which,  like  S03,  are  acid  anhydrides,  and  show  the 
elements  to  be  sexivalent.  They  give  also  acids  of  the  form  H2X04, 
such  as  chromic  acid  H2Cr04.  These  acids  correspond  to  sulphuric 
acid,  and  their  salts,  for  example  the  chromates,  resemble  the 
sulphates. 

Aside  from  the  chromates,  the  first  element  forms  also  two  basic 
hydroxides  Cr(OH)2  and  Cr(OH)3,  from  which  the  numerous  chro- 
mous  (Cr++)  and  chromic  (Cr+++)  salts  are  derived.  Uranium  is 
base-forming,  as  well  as  acid-forming.  Molybdenum  and  tungsten 
are  not  base-forming  elements. 

CHROMIUM  Cr 

The  Chemical  Relations  of  the  Element.  —  Chromium  gives 
four  classes  of  compounds,  and  most  of  them  are  colored  sub- 
stances (Gk.  xP<Va>  color).  The  chromates  are  derived  from 
chromic  acid  H2Cr04,  which,  however,  is  itself  unstable,  and  leaves 
the  anhydride  when  the  solution  is  evaporated.  The  oxide  and 
hydroxide  in  which  the  element  is  trivalent,  namely  Cr2O3  and 
Cr(OH)3,  are  weakly  basic  and  still  more  weakly  acidic.  Hence 
we  have  chromic  salts  such  as  CrCl3  and  Cr2(SO4)3  which  are 
somewhat  hydrolyzed,  but  no  carbonate,  and  no  sulphide  which 
is  stable  in  water.  The  compounds  in  which  the  same  hydroxide 

595 


596  COLLEGE    CHEMISTRY 

acts  as  an  acid  are  the  chromites,  and  are  derived  from  the  less 
completely  hydrated  form  of  the  oxide  CrO(OH).  Potassium 
chromite  K.CrO2  is  more  easily  hydrolyzed,  however,  than  is 
potassium  zincate  or  potassium  aluminate.  Finally,  the  chro- 
mous  salts  such  as  CrCl2  and  CrSC>4  correspond  to  chromous 
hydroxide  Cr(OH)2  in  which  the  element  is  bivalent.  This  hy- 
droxide is  more  distinctly  basic  than  is  chromic  hydroxide,  and 
forms  a  carbonate  and  sulphide  which  can  be  precipitated  in 
aqueous  solution. 

Occurrence  and  Isolation.  —  Chromium  is  found  chiefly  in 
ferrous  chromite  Fe(CrO2)2,  which  constitutes  the  mineral  chro- 
mite, and  in  crocoisite  PbCr04,  which  is  chromate  of  lead.  It 
was  first  discovered  in  the  latter  mineral  by  Vauquelin  (1797). 
The  metal  is  easily  made  by  reduction  of  the  oxide  with  aluminium 
filings  by  Goldschmidt's  method  (p.  556). 

Physical  and  Chemical  Properties.  —  Chromium  is  a  white, 
crystalline,  very  hard  metal  (m.-p.  1520°).  It  does  not  tarnish, 
but  when  heated  it  burns  in  oxygen,  giving  the  green  chromic 
oxide  Cr20s.  It  seems  to  exist  in  two  states,  an  active  and  a  pas- 
sive one,  the  relations  of  which  are  still  somewhat  obscure.  A 
fragment  which  has  been  made  by  the  Goldschmidt  method,  or 
has  been  dipped  in  nitric  acid,  is  passive,  and  does  not  displace 
hydrogen  from  hydrochloric  acid.  When,  however,  the  specimen 
is  warmed  with  this  acid,  it  begins  to  interact,  and  thereafter 
behaves  as  if  it  lay  between  zinc  and  cadmium  in  the  electro- 
motive series.  If  left  in  the  air,  it  slowly  becomes  inactive  again. 

Tin  and  iron  with  hydrochloric  acid  form  stannous  and  ferrous 
chlorides  respectively,  because  the  higher  chlorides,  if  present, 
would  be  reduced  by  the  active  hydrogen  (p.  360).  Here,  for 
the  same  reason,  chromous  chloride  and  not  chromic  chloride  is 
formed : 

Cr  +  2HC1  -»  CrCl2  +  H2,     or    Cr  +  2H+  -> Cr++  +  H2. 

Chromium  is  used  in  making  chrome-steel,  for  armorplate. 
The  strange  alloys,  which,  although  composed  of  active  metals, 
are  not  attacked  by  acids  (even  boiling  nitric  acid),  usually  contain 
chromium  (e.g.,  60%  Cr,  36%  Fe,  4%  Mo). 


THE   CHROMIUM   FAMILY  597 

DERIVATIVE  or  CHROMIC  ACID 

Potassium  Chromate  K2CrO±.  —  This  and  the  sodium  salt,  or 
rather  the  corresponding  dichromates  (see  below),  are  made  di- 
rectly from  chromite,  and  form  the  starting-point  in  the  prepara- 
tion of  the  other  compounds  of  chromium.  The  finely  powdered 
mineral  is  mixed  with  potash  and  limestone,  and  roasted.  The 
lime  is  employed  chiefly  to  keep  the  mass  porous  and  accessible 
to  the  oxygen  of  the  air,  the  potassium  compounds  being  easily 
fusible: 

4Fe(Cr02)2  +  8K2C03  +  702  -»  2Fe203  +  8K2Cr04  +  8C02. 

The  iron  is  oxidized  to  ferric  oxide,  and  the  chromium  passes  from 
the  state  of  chromic  oxide  in  the  chromite  (FeO,Cr203)  to  that  of 
chromic  anhydride  in  the  potassium  chromate  (K2O,CrO3).  Thus, 
more  insight  is  given  into  the  nature  of  the  action  by  the  equation: 

4(FeO,Cr203) +8(K20,C02) +702->2Fe203+8(K20,Cr03) +8C02. 

The  cinder  is  treated  with  hot  potassium  sulphate  solution.  This 
interacts  with  the  calcium  chromate,  which  is  formed  at  the  same 
time,  giving  insoluble  calcium  sulphate: 

CaCr04  +  K2S04  *?  CaS04  j  +  K2Cr04. 

The  whole  of  the  potassium  chromate  goes  into  solution. 

Potassium  chromate  is  pale-yellow  in  color,  gives  anhydrous, 
rhombic  crystals  like  those  of  potassium  sulphate,  and  is  very 
soluble  in  water  (61  :  100  at  10°). 

Sodium  chromate  Na2CrO4,10H20  is  made  by  using  sodium  car- 
bonate in  the  process  just  described. 

The  Dichromates.  —  When  a  solution  of  potassium  sulphate  is 
mixed  with  an  equivalent  amount  of  sulphuric  acid,  potassium 
bisulphate  is  obtainable  by  evaporation:  K2S04  +  H2SO4  — » 
2KHS04.  The  dry  acid  salt,  when  heated,  loses  water  (p.  286), 
giving  the  pyrosulphate  (or  disulphate) :  2KHS04  +±  K2S207  + 
H20,  but  the  latter,  when  redissolved,.  returns  to  the  condition  of 
acid  sulphate.  The  second  action  is  instantly  reversed  in  presence 
of  water.  Now,  when  an  acid  is  added  to  a  chromate  we  should 
expect  the  chromic  acid  H2Cr04,  thus  liberated,  to  interact,  giving 


598  COLLEGE    CHEMISTRY 

an  acid  chromate  (say,  KHCrO4).  No  acid  chromates  are  known, 
however,  and  instead  of  them,  pyrochromates  or  dichromates  are 
produced,  with  elimination  of  water.  In  other  words,  the  second 
of  the  above  actions  is  not  appreciably  reversible  in  presence  of 
water  when  chromates  are  in  question  : 

K2Cr04   +H2S04     -»(H2Cr04)+K2S04. 
K2Cr04  (  +  H2Cr04)  -»  K2Cr207  +  H2O.  _ 
2K2Cr04  +  H2SO4     -»  K2Cr2O7  +  H2O  +  K2SO4.  (1) 

In  terms  of  the  ionic  hypothesis,  S207—  is  unstable  in  water,  and 
interacts  with  the  OH~  ion  it  contains,  giving  water  and  sul- 
phate-ion, while  Cr207—  is  stable  in  water  and  is  formed  from 
the  interaction  of  water  and  chromate-ion  : 


Cr207=  +  20H~  «=$  H2O  +  2Cr04=  (2) 

The  dichromates  of  potassium  and  sodium  are  made  by  adding 
sulphuric  acid  to  the  crude  solution  of  the  chromate  obtained  from 
chromite  (p.  597).  They  crystallize  when  the  liquid  cools,  and  the 
mother-liquor,  containing  the  potassium  sulphate  and  undeposited 
dichromate,  is  used  for  extracting  a  fresh  portion  of  cinder.  As  the 
dichromates  are  much  less  soluble  than  the  chromates,  they  crys- 
tallize from  less  concentrated  solutions,  and  can  therefore  be  ob- 
tained in  purer  condition.  For  this  reason  the  extract  is  always 
treated  for  dichromate. 

Potassium  dichromate  K2Cr207  (or  K2Cr04,Cr03)  crystallizes  in 
asymmetric  tables  of  orange-red  color.  Its  solubility  in  water  is 
8  :  100  at  10°  and  12.5  :  100  at  20°.  Sodium  dichromate  Na^CrgO,, 
2H2O  forms  red  crystals  also,  and  its  solubility  is  109  :  100  at  15°. 
This  salt  is  now  cheaper  than  potassium  dichromate,  and  has 
largely  displaced  the  latter  for  commercial  purposes. 

Chemical  Properties  of  the  Bichromates.  —  1.  When  con- 
centrated sulphuric  acid  is  added  to  a  strong  solution  of  a  dichro- 
mate (or  chromate),  chromic  anhydride  Cr03  separates  in  red 
needles  : 

Na2Cr207  +  H2SO4  ->  Na2S04  +  H20  +  2Cr03  J,  . 

2.  Although  a  dichromate  lacks  the  hydrogen,  it  is  essentially  of 
the  nature  of  an  acid  salt,  just  as  SbOCl  lacks  hydroxyl,  but  is 


THE    CHROMIUM   FAMILY  599 

essentially  a  basic  salt.  Hence,  when  potassium  hydroxide  is 
added  to  a  solution  of  potassium  dichromate,  potassium  chromate 
is  formed  : 

K2Cr2O7  +  2KOH  -»  2K2Cr04  +  H20. 

The  solution  changes  from  red  to  yellow,  and  the  chromate  is 
obtained  by  evaporation.  In  this  way  the  pure  alkali  chromates 
are  made. 

3.  By  addition  of  potassium  dichromate  to  a  solution  of  a  salt  of 
a  metal  whose  chromate  is  insoluble,  the  chromate  and  not  the 
dichromate  is  precipitated.     This  occurs  in  consequence  of  the 
fact  that  there  is  always  a  little  hydrogen-ion  and  CrO4-  (equation 
(2),  above)  in  the  solution  of  the  dichromate: 

2Ba(N03)2  4-  K2Cr207  +  H20  <=>  2BaCr04  J,  +  2KN03  +  2HN03. 

Being  essentially  an  acid  salt,  the  dichromate  produces  a  salt  and 
an  acid,  as  any  acid  salt  would  do.  For  example: 

Ba(N03)2  +  KHS04  <=±  BaSO4  j  +  KN03  +  HN03. 

4.  The  dichromates  of  potassium  and  sodium  melt  when  heated 
and,  at  a  white  heat,  decompose,  giving  the  chromate,  chromic 
oxide,  and  free  oxygen.     To  make  the  equation,  we  note  that  the 
dichromate,  for  example  K2Cr207,  may  be  written  as  K2Cr04,Cr03, 
and  the  CrO3,  if  alone,  will  decompose  thus  :  2Cr03  —  »  Cr2O3  +  30. 
Since  the  product  must  contain  a  multiple  of  02,  the  equation  is: 

4K2Cr207  -*  4K2Cr04  +  2Cr203  +  302. 

5.  With  free  acids  the  dichromates  give  powerful  oxidizing  mix- 
tures, in  consequence  of  their  tendency  to  form  chromic  salts. 
Since  the  former  correspond  to  the  oxide  Cr03  and  the  latter  to 
Cr2O3,  the  passage  from  the  former  to  the  latter  must  furnish  3O 
for  every  2Cr03  transformed.     In  dilute  solutions,  unless  a  body 
capable  of  being  oxidized  is  present,  no  actual  decomposition, 
beyond  the  liberation  of  chromic  acid,*  occurs.     When  concen- 
trated hydrochloric  acid  is  used,  this  acid  itself  suffers  oxidation: 

K2Cr207  +    8HC1  ->  2KC1  +  2CrCl3  +  4H20  (+  30). 
(30)        +    6HC1-+3H2O 


__ 
K2Cr2O7  +  14HC1  ->  2KC1  +  2CrCl3  +  7H2O  +  3C12. 

*  Not  shown  as  a  distinct  stage  in  the  subsequent  equations. 


600  COLLEGE   CHEMISTRY 

When  sulphuric  acid  is  employed,  an  oxidizable  substance  such  as 
hydrogen  sulphide  (cf.  p.  270),  sulphurous  acid,  or  alcohol  must  be 
present,  if  the  dichromate  is  to  be  reduced : 

K2O207  +     4H2S04  -> K2S04  +  Cr2(S04)3  +  4H20(  +  30)     (1) 

(30)  +     3H2S03  -»  3H2S04  (2) 

or      (30)  +  3C2H6OH  ->  3C2H40  |  +  3H2O  (20 

[alcohol]  [aldehyde] 

In  each  case  the  usual  summation  of  (1)  and  (2),  with  omission  of 
the  3O,  gives  the  equation  for  the  whole  action.  When  (1)  is  dis- 
sected, K2O,2CrO3  giving  Cr2O3,3S03  +  3O  is  found  to  be  its  essen- 
tial content.  In  practice,  this  sort  of  action  is  used  for  the  purpose 
of  making  chromic  salts,  and  for  its  oxidizing  effects,  as  in  the 
preparation  of  aldehyde  and  in  the  dichromate  battery. 

Other  Uses  of  Dichromates.  —  When  paper  is  coated  with 
gelatine  containing  a  soluble  chromate  or  dichromate  and,  after 
being  dried,  is  exposed  to  light,  chromic  oxide  is  formed  by  reduc- 
tion, and  combines  with  the  gelatine.  This  product  will  not  swell 
up  or  dissolve  in  tepid  water,  as  does  pure  gelatine.  This  action 
is  used  in  many  ways  for  purposes  of  artistic  reproduction.  Thus, 
if  the  gelatine  mixture  is  made  up  with  lampblack  and,  after  the 
coating  has  dried,  is  covered  with  a  negative  and  exposed  to  light, 
the  parts  which  were  protected  from  illumination  may  afterwards 
be  washed  away,  while  the  carbon  print  remains.  The  gelatine 
layer  can  be  transferred  to  wood  or  copper  before  washing.  When 
materials  of  different  colors  are  substituted  for  the  lampblack, 
prints  of  any  desired  tint  may  be  made  by  the  same  process. 

Sodium  dichromate  is  used,  instead  of  tan-bark,  in  tanning  kid 
and  glove  leathers.  A  reducing  agent  is  employed  to  precipitate 
chromic  hydroxide  Cr(OH)3  in  the  leather.  Its  use  diminishes 
the  time  required  for  the  process  from  8  or  10  months  to  a  few 
hours.  The  hide  is  a  mixture  of  colloidal  materials,  and  the  hy- 
droxide is  adsorbed. 

Insoluble  Chromates.  —  A  number  of  chromates,  formed  by 
precipitation  with  a  solution  of  a  soluble  chromate  or  dichromate, 
are  familiar.  Thus,  lead  chromate  PbCr04  is  used  as  a  yellow 
pigment.  By  treatment  with  limewater  it  gives  a  basic  salt  of 
brilliant  orange  color  —  chrome-red  Pb2OCr04.  Salts  of  calcium 


THE    CHROMIUM   FAMILY  601 

give  a  yellow,  hydrated  calcium  chromate  CaCr04,2HO2  analogous 
to  gypsum,  an'd,  like  it,  perceptibly  soluble  in  water  (0.4  :  100  at 
14°) .  Barium  chromate  BaCr04  is  also  yellow.  It  interacts  with 
active  acids  to  form  the  dichromate,  and  passes  into  solution.  It 
is  not  soluble  enough  to  be  attacked  by  acetic  acid.  Strontium 
chromate  SrCr04,  however,  is  soluble  in  acetic  acid.  Silver  chro- 
mate Ag2CrO4  is  red,  and  interacts  easily  with  acids.  It  will  be 
observed  that  there  is  a  close  correspondence  between  the  relative 
solubilities  (see  Table)  of  the  chromates  and  the  sulphates. 

Chromic  Anhydride  CrO3.  —  This  oxide  is  made  as  described 
above  (par.  1,  p.  598),  and  is  often  called  chromic  acid.  It  is 
soluble  in  water,  and  combines  with  the  latter  to  some  extent, 
giving  dichromic  acid  H2.Cr207.  In  a  solution  acidified  with  an 
active  acid  it  is  much  used  as  an  oxidizing  agent  for  organic  sub- 
stances. It  interacts  with  acids  in  the  same  way  as  do  the  dichro- 
mates,  giving  chromic  salts  and  furnishing  oxygen  to  the  oxidizable 
body.  When  heated  by  itself,  it  loses  oxygen  readily,  and  yields 
the  green  chromic  oxide :  4Cr03  — •>  2Cr2O3  +  302. 

Chromyl  Chloride  OO2O3.  —  This  compound  corresponds  to 
sulphuryl  chloride  S02C12,  and  is  made  by  distilling  a  dichromate 
with  a  chloride  and  concentrated  sulphuric  acid: 

K2Cr207  +  4KC1  +  3H2SO4  ->  2Cr02Cl2  t  +  3K2SO4  +  3H2O. 

The  hydrochloric  acid  liberated  from  the  chloride  may  be  supposed 
to  interact  with  chromic  acid  from  the  dichromate: 

CrO2(OH)2  +  2HC1  ~>  CrO2Cl2  +  2H2O. 

Chromyl  chloride  is  a  red  liquid,  boiling  at  1 18°.  It  fumes  strongly 
in  moist  air,  being  hydrolyzed  by  water.  This  action  is  the  re- 
verse of  that  shown  in  the  last  equation.  The  corresponding 
bromine  and  iodine  compounds  are  unstable,  and  when  a  bromide 
or  iodide  is  treated  as  described  above,  the  halogens  are  liberated 
by  oxidation,  and  no  volatile  compound  of  chromium  appears. 
Hence,  when  an  unknown  halide  is  mixed  with  potassium  dichro- 
mate and  sulphuric  acid,  and  distilled,  and  the  vapors  are  caught 
in  ammonium  hydroxide,  the  finding  of  a  chromate  in  the  dis- 


602  COLLEGE    CHEMISTRY 

tillate  demonstrates  the  existence  of  a  chloride  in  the  original 
substance : 

Cr02Cl2  +  4NH4OH  ->  (NH4)2CrO4  +  2NH4C1  +  2H2O. 

This  action  is  used  as  a  test  for  the  presence  of  traces  of  chlorides 
in  large  amounts  of  bromides  or  iodides. 

CHROMIC  AND  CHROMOUS  COMPOUNDS 

Chromic  Chloride.  —  A  hydrated  chloride  CrCl3,6H2O  is  ob- 
tained by  treating  the  hydroxide  Cr(OH)3  with  hydrochloric  acid 
and  evaporating.  When  heated,  this  hydrate  is  hydrolyzed,  and 
chromic  oxide  remains.  The  anhydrous  chloride  CrCl3  is  formed 
by  sublimation,  as  a  mass  of  brilliant,  reddish-violet  scales,  when 
chlorine  is  led  over  heated  metallic  chromium.  In  this  form  the 
substance  dissolves  with  extreme  slowness,  even  in  boiling  water, 
but  in  presence  of  a  trace  of  chromous  chloride  or  stannous  chloride 
it  is  easily  soluble.  The  solution  is  green,  as  are  all  solutions  of 
chromic  salts  after  they  have  been  boiled,  but  on  standing  in  the 
cold,  bluish  crystals  of  CrCl3,6H20  are  deposited.  These  give  a 
violet  solution  containing  Cr+++  +  3C1~,  but  boiling  reproduces  the 
green  color.  The  green  material  can  also  be  obtained  in  crystals 
as  a  hexahydrate,  and  is  therefore  isomeric  (p.  421)  with  the 
violet  variety.  With  the  green  isomer,  in  cold  solution,  silver 
nitrate  precipitates  at  first  only  one-third  of  the  chlorine  as  silver 
chloride. 

Chromic  Hydroxide.  —  When  ammonium  hydroxide  is  added 
to  a  solution  of  a  chromic  salt,  a  hydrated  hydroxide  of  pale-blue 
color,  2Cr(OH)3,H20,  is  thrown  down.  This  interacts  with  acids, 
giving  chromic  salts.  It  also  dissolves  in  potassium  and  sodium 
hydroxides  to  form  green  solutions  of  chromites  of  the  form  KCr02. 
When  the  solutions  of  the  alkali  chromites  are  boiled,  the  free 
chromic  hydroxide,  present  in  consequence  of  hydrolysis,  is  con- 
verted into  a  greenish,  less  completely  hydrated,  and  less  soluble 
variety.  This  begins  to  come  out  as  a  precipitate,  and  soon  the 
whole  action  is  reversed.  Insoluble  chromites,  such  as  that  of 
iron  Fe(CrO2)2,  are  found  in  nature.  Many  of  them,  like  Zn(Cr02)2 
and  Mg(Cr02)2,  may  be  formed  by  fusing  the  oxide  of  the  metal 


THE    CHROMIUM   FAMILY  603 

with  chromic  oxide;  the  action  being  similar  to  that  used  in  making 
zincates  (p.  529)  and  aluminates  (p.  557).  The  hydroxide  is  used 
as  a  mordant  (p.  565)  and  is  the  active  substance  in  the  chrome- 
tanning  process  (p.  600). 

Chromic  Oxide  O2O3.  —  This  oxide  is  obtained  as  a  green, 
infusible  powder  by  heating  the  hydroxide;  or,  more  readily,  by 
heating  dry  ammonium  dichromate;  or  by  igniting  potassium 
dichromate  with  sulphur  and  washing  the  potassium  sulphate  out 
of  the  residue : 

(NH4)2Cr207  ->  N2  +  4H20  +  Cr2O3, 

K2Cr207  +  S  ->  K2S04  +  Cr203. 

Chromic  oxide  is  not  affected  by  acids,  but  may  be  converted  into 
the  sulphate  by  fusion  with  potassium  bisulphate.  It  is  used  for 
making  green  paint,  and  for  giving  a  green  tint  to  glass.  When  the 
oxide,  or  any  of  the  chromic  salts,  is  fused  with  a  basic  substance 
such  as  an  alkali  carbonate,  it  passes  into  the  form  of  a  chromate, 
absorbing  the  necessary  oxygen  from  the  air.  If  an  alkali  nitrate 
or  chlorate  is  added,  the  oxidation  goes  on  more  quickly.  The 
alkaline  solution  of  the  chromites  may  be  oxidized,  for  example  by 
adding  chlorine  or  bromine,  and  chromates  are  formed. 

Chromic  Sulphate  Cr8(SO4)3,15/f3O.  —  This  salt  crystallizes 
in  reddish-violet  crystals,  and  may  be  made  by  treating  the  hy- 
droxide with  sulphuric  acid.  When  mixed  with  potassium  sul- 
phate, it  gives  reddish-violet,  octahedral  crystals  of  chrome-alum 
(cf.  p.  558),  K2SO4,Cr2(S04)3,24H2O.  This  double  salt  is  most 
easily  obtained  by  reducing  potassium  dichromate  in  dilute  sul- 
phuric acid  by  means  of  sulphurous  acid  (p.  600),  and  allowing 
the  solution  to  crystallize.  The  solution  of  the  crystals,  either  of 
the  pure  sulphate  or  of  the  alum,  is  bluish-violet  (Cr"H~+),  but 
when  boiled  becomes  green.  The  green  compound  is  formed  by 
hydrolysis  and  is  gummy  and  uncrystallizable.  It  even  yields 
products  which  do  not  show  the  presence  either  of  the  Cr4"^  or 
the  SO4—  ion.  It  seems  to  be  formed  thus: 

2Cr2(S04)3  +  H20  ?=»  Cr40(S04)4.S04  +  H2SO4. 

The  green  materials  revert  slowly  to  the  violet  ones  by  reversal  of 
the  above  action  when  the  solution  remains  in  the  cold,  and  so 


604  COLLEGE    CHEMISTRY 

crystals  of  the  sulphate  or  of  the  alum  are  obtainable  from  the 
green  solutions. 

Chromous  Compounds.  —  By  the  interaction  of  chromium 
with  hydrochloric  acid,  or  by  reducing  chromic  chloride  in  a  stream 
of  hydrogen,  chromous  chloride  CrCl2  is  formed.  The  anhydrous 
salt  is  colorless,  and  its  solution  is  light  blue  (Cr++).  Like  stan- 
nous  chloride,  it  is  very  easily  oxidized  by  the  air,  a  solution  of  it 
containing  excess  of  hydrochloric  acid  being  used  in  the  laboratory 
to  absorb  oxygen: 

4CrCl2  +  4HC1  +  O2  -+4CrCl3  +  2H2O. 

Chromous  hydroxide  Cr(OH)2  is  obtained  as  a  yellow  precipitate 
when  alkalies  are  added  to  the  chloride.  With  sulphuric  acid  it 
gives  chromous  sulphate  CrSO4,7H20,  which  is  one  of  the  vitriols 
(p.  529). 

Analytical  Reactions  of  Chromium  Compounds.  —  The 

chromic  salts  give  the  bluish-violet  chromic-ion  0+++,  or  the  green 
complex  cations,  and  may  be  recognized  in  solution  by  their  color. 
The  chromates  and  dichromates  give  the  ions  CrO4—  and  Cr207~, 
which  are  yellow  and  red  respectively.  From  chromic  salts, 
alkalies  and  ammonium  sulphide  precipitate  the  bluish-green 
hydroxide,  and  carbonates  give  a  basic  carbonate  which  is  almost 
completely  hydrolyzed  to  hydroxide.  By  fusion  with  sodium 
carbonate  and  sodium  nitrate,  they  yield  a  yellow  bead  containing 
the  chromate.  The  chromates  and  dichromates  are  recognized 
by  the  insoluble  chromates  which  they  precipitate,  and  by  their 
oxidizing  power  when  mixed  with  acids.  All  compounds  of  chro- 
mium give  a  green  borax  bead  containing  chromic  borate,  and  this 
bead  differs  from  that  given  by  compounds  of  copper  (cf.  p.  510), 
both  in  tint  and  in  being  unreducible. 

MOLYBDENUM,  TUNGSTEN,  URANIUM 

Molybdenum.  —  This  element  is  found  chiefly  in  wulfenite 
PbMoO4  and  molybdenite  MoS2.  The  latter  resembles  black 
lead  (graphite),  and  its  appearance  suggested  the  name  of  the 
element  (Gk.  noXvfiSaiva,  lead).  The  molybdenite  is  converted  by 


THE   CHROMIUM   FAMILY  605 

roasting  into  molybdic  anhydride  MoO3.  When  this  is  treated  with 
ammonium  hydroxide,  or  with  sodium  hydroxide,  ammonium 
molybdate  (NH4)2Mo04  or  sodium  molybdate  Na2Mo04,10H20  is 
obtained.  The  metal  itself  is  liberated  by  reducing  the  oxide  or 
chloride  with  hydrogen.  '  When  pure  it  is  a  silvery  metal  and, 
like  iron  (q.v.),  takes  up  carbon  and  shows  the  phenomena  of 
tempering.  The  oxides  Mo2O3,  MoO2,  and  MoO3  are  known,  but 
the  lower  oxides  are  not  basic.  The  chlorides  Mo3Cl6,  MoCl3, 
MoCLi,  and  MoCl5  have  been  made.  The  chief  use  of  molybdenum 
compounds  in  the  laboratory  is  in  testing  for  and  estimating  phos- 
phoric acid.  When  a  little  of  a  phosphate  is  added  to  a  solution 
of  ammonium  molybdate  in  nitric  acid,  and  the  mixture  is  warmed, 
a  copious  yellow  precipitate  of  a  phosphomolybdate  of  ammonium 
(NH4)3P04,llMoO3,6H2O  is  formed.  The  compound  is  soluble 
in  excess  of  phosphoric  acid  and  in  alkalies,  but  not  in  dilute 
mineral  acids. 

Tungsten.  —  The  minerals  scheelite  CaW04  and  wolfram 
[Fe,Mn]W04  are  tungstates  of  calcium  and  of  iron  and  manganese, 
respectively.  By  fusion  of  wolfram  with  sodium  carbonate  and 
extraction  with  water,  sodium  tungstate  Na2W04,2H20  is  secured. 
It  is  used  as  a  mordant  and  for  rendering  muslin  fireproof.  Acids 
precipitate  tungstic  acid  H2W04,H20  from  solutions  of  this  salt. 
The  element  gives  the  oxides  WO2  and  W03,  the  latter  being 
formed  by  ignition  of  tungstic  acid.  The  chlorides  WC12,  WC14, 
WCU,  and  WC16  are  known,  the  last  being  formed  directly,  and 
the  others  by  reduction. 

The  metal  has  important  uses,  and  the  annual  production  is 
greater  than  the  total  of  all  the  metals  which  follow  it  in  the  list 
on  p.  436.  The  metal  (sp.  gr.  19.6)  can  be  liberated  by  reduction 
of  the  oxide  by  hydrogen  or  by  carbon.  It  has  a  higher  melting 
point  (3540°)  than  any  other  metal  and,  on  this  account,  and  be- 
cause it  is  less  volatile  than  carbon,  is  now  used  for  filaments  in 
electric  lamps.  A  carbon  filament  also  requires  3.25  watts  per 
candle  power  while  a  tungsten  filament  uses  only  1.25  watts  per 
1  c.  p.  The  powdered  metal  obtained  by  reduction  can  be  pressed 
into  wire  form  and  then  rolled  while  strongly  heated  by  an  electric 
current  until  a  compact  wire  is  obtained.  The  metal  can  also  be 
obtained  in  massive  form  by  reducing  the  oxide  with  aluminium, 


606  COLLEGE    CHEMISTRY 

provided  the  crucible  and  mixture  are  heated  strongly  in  advance. 
In  1914,  in  the  United  States  alone,  about  a  hundred  million  tung- 
sten lamps  were  manufactured.  Shop  work  has  been  almost  revolu- 
tionized by  the- use  of  tungsten  steel  tools,  which  can  be  used  at 
high  speed  and,  even  when  thus  heated  red  hot  by  friction,  retain 
their  temper.  Tungsten  steel  contains  tungsten  (16  to  20%), 
carbon  (0.55  to  0.75%),  chromium  (2.5  to  5%),  and  vanadium 
(0.35  to  1.5%). 

Uranium.  —  Pitchblende,  which  contains  the  oxide  U3O8  along 
with  smaller  amounts  of  many  other  elements,  is  found  mainly 
in  Joachimsthal  (Bohemia)  and  in  Cornwall.  Carnotite,  a  ura- 
nate  and  vanadate  of  potassium  K20,2UO3,V205,3H2O  occurs  in 
Colorado.  Pitchblende  is  roasted  with  lime,  the  calcium  uranate 
CaU04  thus  formed  is  decomposed  with  sulphuric  acid,  giving 
uranyl  sulphate  UO2SO^  When  excess  of  sodium  carbonate  is 
added  to  the  solution  of  the  latter,  the  foreign  metals  are  precipi- 
tated and  sodium  diuranate  Na^C^TI^O,  which  is  also  thrown 
down,  dissolves  in  the  excess  as  Na^UC^. 

After  filtration,  the  diuranate  of  sodium  is  reprecipitated  by 
neutralizing  with  sulphuric  acid  and  boiling.  This  salt  is  used 
in  making  uranium  glass,  which  shows  a  yellowish-green  fluores- 
cence. The  property  is  due  to  the  fact  that  the  wave-lengths  of 
part  of  the  invisible,  ultra-violet  rays  of  the  sunlight  are  lengthened, 
and  a  greenish  light  is  therefore  in  excess.  The  oxides  are  UO2  a 
basic  oxide,  U203,  U308  the  most  stable  oxide,  U03  uranic  anhy- 
dride, and  UC>4  a  peroxide. 

When  the  oxide  U02  is  treated  with  acids,  it  gives  uranous  salts 
such  as  uranous  sulphate  U(S04)2,4H20.  Uranic  anhydride  and 
uranic  acid  interact  with  acids,  giving  basic  salts,  such  as  U02SO4, 
3|H2O,  and  U02(NO3)2,6H20,  which  are  named  uranyl  sulphate, 
uranyl  nitrate,  and  so  forth.  They  are  yellow  in  color,  with  green 
fluorescence.  Ammonium  sulphide  throws  down  the  brown,  un- 
stable uranyl  sulphide  UO2S  from  their  solutions. 

RADIOACTIVE  ELEMENTS 

Historical.  —  We  have  seen  (p.  303)  that  in  an  evacuated  tube, 
through  which  an  electric  discharge  is  passed,  the  "rays"  emanat- 
ing from  the  cathode  (cathode  rays)  strike  the  anti-cathode  and 


THE   RADIOACTIVE    ELEMENTS  607 

the  glass  beyond  it.  They  produce  in  the  glass  a  greenish-yellow, 
fluorescent  light.  These  "rays"  were  discovered  by  Sir  William 
Crookes  (1878),  and  later  were  shown  to  consist  of  particles  of 
negative  electricity  or  electrons,  each  having  a  mass  about  T?Viy  of 
an  atom  of  hydrogen.  Rontgen  (1895)  accidentally  discovered 
that  the  fluorescent  light  (X-rays)  could  penetrate  paper,  flesh, 
and  other  materials  composed  of  elements  of  low  atomic  weight 
and  acted  upon  photographic  plates.  In  1896  Henri  Becquerel 
observed  that  minerals  containing  uranium  gave  off  a  sort  of 
radiation  which  could  penetrate  black  paper  that  was  opaque  to 
ordinary  light  and  reduce  the  silver  bromide  on  a  photographic 
plate  placed  beneath  the  paper.  He  also  discovered  that  an 
electrometer  (Fig.  133),  in  which  the  gold  leaves  had  been  caused 
to  separate  by  charging  with  electricity, 
lost  its  charge  rapidly  when  the  uranium 
ore  (or  salt)  was  brought  near  (3-4  cm.) 
to  the  knob  connected  with  the  leaves. 
The  uranium  material  rendered  the  air 
a  conductor  (" ionized"  the  air)  and  this 
effect  permitted  the  escape  of  the  electric 
charge,  which  otherwise  would  have  been 
retained  for  a  considerable  time.  In  the 
quantitative  measurement  of  radioactiv-  FIQ 

ity,  we  now  compare  the  times  required 

for  the  discharge  of  an  electroscope  by  different  specimens  of  radio- 
active matter.  The  presence  of  10~12  g.  of  such  matter  can  thus 
be  detected. 

The  radioactivity  of  every  pure  uranium  compound  is  propor- 
tional to  its  uranium  content.  The  ores  are,  however,  relatively 
four  times  as  active.  This  fact  led  M.  and  Mme.  Curie,  just  after 
1896,  to  the  discovery  that  the  pitchblende  residues,  from  which 
practically  all  of  the  uranium  had  been  extracted,  were  neverthe- 
less quite  active.  About  a  ton  of  the  very  complex  residues 
having  been  separated  laboriously  into  the  components,  it  was 
found  that  a  large  part  of  the  radioactivity  remained  with  the 
sulphate  of  barium.  From  this  barium  sulphate,  a  product  free 
from  barium,  and  at  least  one  million  times  more  active  than 
uranium,  was  finally  secured  in  the  form  of  the  bromide.  The 
nature  of  the  spectrum  and  the  chemical  relations  of  the  element, 


608  COLLEGE    CHEMISTRY 

now  named  radium,  placed  it  with  the  metals  of  the  alkaline 
earths.  The  ratio  by  weight  of  chlorine  to  radium  in  the  chloride 
is  35.46  :  113,  so  that,  on  the  assumption  that  the  element  is  bi- 
valent, its  chloride  is  RaCl2  and  its  atomic  weight  is  226.  With 
this  value  it  occupies  a  place  formerly  vacant  in  the  periodic  table. 
In  1910  Mme.  Curie  obtained  metallic  radium  by  electrolyzing 
a  solution  of  radium  chloride,  using  a  mercury  cathode,  and  ex- 
pelling the  mercury  by  distillation.  It  was  a  white  metal  (m.-p. 
700°)  which,  like  calcium,  quickly  tarnished  in  the  air  and  dis- 
placed hydrogen  from  water. 

The  Nature  of  the  "Rays."  —Many  properties  show  that 
the  "rays"  emitted  by  compounds  of  uranium  and  of  radium  are 
of  three  kinds.  They  are  most  sharply  distinguished  from  one 
another  when  allowed  to  pass  through  a  powerful  magnetic  field. 
The  alpha-rays  are  positively  charged  and  are  bent  in  one  direction 
while  the  beta-rays  are  negative  and  are  bent  in  the  other.  The 
gamma-rays  are  not  affected. 

The  alpha-rays  are  atoms  of  helium  (p.  336)  thrown  off  in 
straight  lines  with  varying  initial  velocities,  averaging  about  one- 
tenth  that  of  light  (say,  30,000  kilometers  per  second.  The 
a-particles  from  Ra-C,  e.g.,  19,220  kilom.  per  sec.).  Each  such 
atom  bears  a  double  positive  charge  (the  unit  being  the  charge  on 
a  univalent  positive  ion),  and  a  delicate  electroscope  readily  in- 
dicates the  entrance  of  a  single  atom.  These  alpha-particles, 
being  each  four  times  as  heavy  as  an  atom  of  hydrogen,  plough 
their  way  through  tens  of  thousands  of  air-molecules  and  usually 
go  about  3-8  cm.  before  being  stopped.  The 
emission  of  atoms  of  helium  can  be  detected  by 
means  of  Crookes  spintharoscope  (Fig.  134). 
The  particle  of  radium  bromide  is  at  B,  and 


some  of  the  charged  helium  atoms  strike  a  sur- 
face C  covered  with  zinc  sulphide,  producing  fault  flashes  of  light. 
The  lens  A  magnifies  the  flashes  and  the  latter  can  be  seen  in  a 
dark  room  after  the  eye  has  become  thoroughly  rested  (15-20 
minutes).  The  helium  gas  given  off  by  radium  compounds  was 
collected  by  Soddy  working  with  Ramsay  and  identified,  and  its 
rate  of  production  was  measured.  The  amount  was  equal  to  158 
cubic  mm.  per  1  g.  of  radium  per  year. 


THE   RADIOACTIVE   ELEMENTS  609 


The  alpha-particles,  in  passing  through  the  air-molecules, 
ionize  the  air,  and  the  ionized  air  has  the  same  power  that  dust 
possesses  (p.  333)  of  affording  nuclei  on  which  moisture  can  con- 
dense. Hence,  when  a  particle  of  a  radium  compound  is  supported 
in  a  flask  containing  air  saturated  with  moisture,  and  the  air  is 
suddenly  cooled  by  expansion,  the  paths  of  the  particles  become 
lines  of  fog.  With  powerful  illumination,  the  fog-tracks  (Fig.  135) 
can  be  photographed  (Wilson),  and  the  lengths  of  the  paths  can 
be  measured. 

The  beta-particles  are  electrons  (p.  303),  or  unit  charges  of  nega- 
tive electricity,  and  are  shot  out  with  a  velocity  approaching  that 
of  light  (300,000  kiloms.  per  sec.).  They  are  therefore  identical 
with  cathode  rays,  but  move  many  times  more  rapidly.  Being  very 
light  (weight,  TsW  of  an  atom  of  hydrogen),  their  paths,  although 
straight  at  first,  soon  become  tortuous  owing  to  collisions  with 
the  relatively  massive  air-molecules.  Half  of  them  are  lost  after 
going  about  4  cm;  Their  fog  tracks  are  fainter  than  are  those  of 
the  a-particles  and  extremely  tangled.  Being  much  lighter  than 
a-particles,  their  paths  are  actually  coiled  into  circles  or  spirals 
by  a  magnetic  field. 

The  gamma-rays  are  identical  with  X-rays  (vibrations  in  the 
ether  of  short  wave-length,  p.  303),  and  are  presumably  produced 
like  the  latter  by  the  impacts  of  the  electrons  on  the  surrounding 
matter. 

The  helium  atoms  are  almost  all  stopped  by  a  sheet  of  paper  or 
by  aluminium  foil  0.1  mm.  thick.  The  electrons  have  greater 
penetrating  power,  many  passing  through  gold-leaf,  but  being 
practically  all  destroyed  by  a  sheet  of  aluminium  1  cm.  thick. 
The  gamma-rays  (X-rays),  however,  are  able  to  penetrate  rela- 
tively thick  layers  of  metals  and  other  materials  of  low  atomic 
weight. 

One  of  the  most  striking  facts  is  that  the  stoppage  by  the  air 
of  so  many  rapidly  moving  particles  results  in  the  production  of 
much  heat.  One  gram  of  radium  would  produce  about  120  cal. 
per  hour. 

Disintegration.  —  The  emission  of  atoms  of  helium  and  of 
electrons  was  first  explained  by  Rutherford  (1902-3),  then  of 
McGill  University,  Montreal,  as  being  due  to  the  spontaneous 


610 


COLLEGE    CHEMISTRY 


FOG-TRACKS  FROM  RADIUM   (C.  T.  R.  WILSON) 

1,  2.  Paths  of  helium  atoms.  3.  Part  of  2,  enlarged.  4.  Paths  of  electrons. 

FIG.  135. 


THE   RADIOACTIVE   ELEMENTS  611 

disintegration  of  the  atoms  of  uranium,  radium,  and  other  radio- 
active elements.  Thus,  Rutherford  was  the  first  to  show  that 
radium  compounds  produced  a  gaseous  substance  called  the 
radium  emanation  (niton),  which  was  the  residue  left  after  the 
emission  of  one  atom  of  helium  from  an  atom  of  radium.  This 
gas  was  itself  radioactive  and  underwent  further  disintegration, 
depositing  a  solid  radioactive  residue  on  bodies  in  contact  with 
it.  Furthermore,  every  known  uranium  ore  contains  radium 
(McCoy)  and  radium  emanation  (Boltwood)  in  amounts  propor- 
tional to  the  uranium  content.  Also,  after  the  radium  has  been 
removed,  the  pure  uranium  compound  gives  off  at  first  only 
a-particles,  but  gradually  recovers  its  whole  radioactivity  and 
is  then  found  to  contain  radium  emanation  once  more  (Soddy). 
It  thus  appears  that  uranium  is  the  starting  point,  and  that  the 
disintegration  proceeds  by  steps,  producing  a  number  of  different 
products.  Each  of  these  is  formed  from  one  such  product  and  by 
disintegration  furnishes  another. 

Unlike  ordinary  chemical  change,  the  rate  of  disintegration  is 
not  affected  by  conditions.  It  can  neither  be  started  nor  stopped 
at  will.  It  is  no  more  vigorous  at  2000°  than  at  -200°.  Other 
changes  occur  between  atoms,  these  within  each  atom. 

The  law,  due  also  to  Rutherford,  describing  the  rate  at  which 
any  one  radioactive  element  disintegrates  is  simple.  Only  a 
certain  fraction  of  the  whole  of  any  one  specimen  undergoes  the 
change  in  unit  time.  Thus,  as  the  total  amount  diminishes  be- 
cause of  the  change,  the  amount  changing  during  the  next  unit 
of  time,  being  a  constant  fraction  of  the  whole,  must  be  less. 
Hence  an  infinite  time  would  be  required  for  the  complete  disin- 
tegration of  any  one  specimen.  For  convenience,  therefore,  it  is 
sometimes  the  custom  to  give  as  a  specific  property  of  each  radio- 
active element  the  time  required  for  the  decay  of  half  its  amount 
and  therefore  the  loss  of  half  of  its  radioactivity.  More  usually, 
the  property  given  is  the  one  called  the  average  life  of  the  element. 
The  value  of  this  is  equal  to  the  inverse  of  the  fraction  disinte- 
grating per  unit  time,  and  is  about  1.44  times  the  period  of  half 
change.  Numerically  it  is  the  sum  of  the  separate  periods  of 
future  existence  of  all  the  atoms  divided  by  the  number  of  such 
atoms  present  at  the  starting  point. 

Radium  emits  helium  atoms  at  the  rate  of  3.4  X  1010  per  gram 


612  COLLEGE   CHEMISTRY 

per  second.  From  this  fact,  we  can  calculate  its  average  life  to 
be  about  2400  years.  Hence,  if  it  were  not  continuously  being 
produced  (from  uranium),  the  whole  supply  would  have  been 
exhausted  long  before  the  earth  reached  a  habitable  condition. 

The  Uranium  Group  of  Radioactive  Elements.  —  The  fol- 
lowing shows  the  various  elements  produced  from  uranium  by 
successive  disintegrations.  When  a  helium  atom  or  an  electron  is 
expelled,  the  fact  is  shown  by  the  symbols  He  and  e,  respectively. 
The  first  number  below  each  element  is  the  average  life  of  that 
member  of  the  series  (y  =  year,  d  =  day,  h  =  hour,  m  =  minute, 
s  =  second).  The  second  number  is  the  atomic  weight,  obtained 
by  subtracting  from  the  at.  wt.  of  uranium  (238.2)  the  weight  (4) 
of  each  helium  atom  emitted. 


i     -»e+U-X2->e  +U2  ->He+Ionium 

8X10"  y.  35.5  d.  1.65m.  3X10«  y.  2X10*  y. 

238.2  234.2  234.2  234.2  230.2 

—  »He+Ra    —  >He+Niton  ->He+Ra-A         -»He+Ra-B 

2440  y.  5.55  d.  4.3m.  38.5m. 

226  222  218  214 

->€  +Ra-C-»€    +Ra-d  -»He+Ra-D        ->€     +Ra-E 

28.1  m.        10-«  s.  24  y.  7.2  d. 

214          214  210  210 

-»e  +Ra-F->He+Pb(end) 

198  d. 

210          206 

A  purified  salt  of  uranium  recovers  half  its  activity  in  about 
three  weeks,  and  reaches  full  equilibrium  in  from  six  months  to 
a  year.  An  equilibrium  is  attained  when  the  speed  at  which  each 
disintegration  product  is  being  formed  is  balanced  by  the  equal 
speed  with  which  it  is  passing  into  the  next  member  of  the  series. 
The  complex  operations  required  for  studying  all  the  members 
of  the  series  cannot  be  given  here.  It  may  be  said,  however, 
that  a  pure  uranium  salt  in  solution  gives  with  ammonium  car- 
bonate a  precipitate  which  is  wholly  soluble  in  excess  of  the  re- 
agent. After  about  a  year,  another  portion  of  the  same  specimen 
leaves  a  slight  precipitate  which  is  insoluble  in  excess  and  contains 
the  products  of  disintegration,  chiefly  U-X  which  was  first  obtained 
in  this  way  by  Crookes. 


THE   RADIOACTIVE    ELEMENTS  613 

The  radium  emanation  was  shown  by  Ramsay  to  be  one  of  the 
inert  gases  (p.  337),  and  was  renamed  niton.  Its  density  was 
determined  .experimentally  with  a  small  sample,  using  a  micro- 
balance  capable  of  weighing  to  1/500,000  mgm.,  and  found  to  be 
222.4  (density  of  oxygen  =  32).  . 

The  end-product  of  the  disintegration  is  lead,  and  all  uranium 
ores  contain  lead.  Lead  from  other  sources  gives  a  chloride 
PbCl2  in  which  207.20  parts  of  lead  are  combined  with  2  X  35.46 
parts  of  chlorine.  The  atomic  weight  207.2  cannot,  however,  be 
reached  by  subtracting  a  whole  number  of  atomic  weights  of 
helium  from  the  atomic  weight  of  uranium,  the  number  206  being 
obtained  instead.  Recently,  lead  chloride  prepared  from  the  lead 
found  in  various  ores  of  uranium  has  been  analyzed  by  Richards 
of  Harvard,  as  well  as,  independently,  by  two  other  chemists, 
and  the  atomic  weight  of  this  lead  was  found  to  be  from  206.4  to 
206.8  in  different  samples.  This  lead  chloride  has  properties 
identical  with  those  of  ordinary  lead  chloride  and  is,  therefore,  by 
definition,  the  same  substance.  Hence  these  investigations  have 
revealed  the  first  known  exception  to  the  law  of  definite  pro- 
portions. 

Since  the  initial  (U)  and  final  (Pb)  materials  are  both  electri- 
cally neutral,  it  must  be  assumed  that  at  some  stages  more  than 
one  electron  per  atom  is  expelled.  8He++  are  lost  and  therefore 
16e~. 

Additional  Data. — The  yield  of  radium  is  very  small.  6000  kg. 
of  pitchblende,  after  extraction  of  the  uranium,  yield  about 
2000  kg.  of  residue.  This  affords  about  6  to  8  kg.  of  the  mixture 
of  radium  and  barium  sulphates,  from  which  0.2  g.  of  pure  radium 
bromide  can  be  prepared. 

One  gram  of  uranium,  after  it  has  produced  the  equilibrium 
proportion  of  radium  (about  3.2  X  10~7  g.),  gives  off  helium  at 
the  rate  of  1  c.c.  in  sixteen  million  years.  Since  the  mineral 
fergusonite  contains  26  c.c.  of  accumulated  helium  for  every 
gram  of  uranium,  the  samples  of  this  mineral  must  be  at  least 
416  million  years  old. 

The  complete  disintegration  of  1  c.c.  of  niton  to  lead  would 
deliver  about  seven  million  calories,  but,  of  course,  the  liberation 
of  the  heat  would  be  spread  over  a  great  length  of  time. 


614  COLLEGE    CHEMISTRY 

Chemical  Actions  of  the  "  Rays."  —  The  radiations  which 
are  most  active  in  ionizing  air  and  in  acting  upon  photographic 
plates  are  the  a-particles.  These  particles  also  cause  the  flashes 
of  light  when  they  encounter  zinc  sulphide.  The  radiations 
change  the  colors  of  minerals,  including  gems,  and  give  a  deep 
violet  color  to  the  glass  tube  containing  the  specimen.  They 
also  turn  atmospheric  oxygen  in  part  into  ozone  and,  in  solution, 
produce  traces  of  hydrogen  peroxide  in  the  water. 

The  radiations  also  destroy  minute  organisms  and  kill  the  cells 
of  the  skin,  producing  sores.  They  have  been  employed  in  the 
treatment  of  lupus  and  of  superficial  cancerous  growths. 

Other  Radioactive  Series.  —  Thorium,  found  as  phosphate 
in  monazite  sand,  is  also  radioactive  and  furnishes  a  series  of 
disintegration  products.  The  final  material  is  a  salt  of  lead. 
Analysis  of  the  chloride  of  lead  made  from  traces  of  the  element 
found  in  all  thorium  minerals  shows  that  the  atomic  weight  (Soddy) 
is  208.4,  while  that  of  ordinary  lead  is  207.2.  The  atom  of  thorium 
(at.  wt.  232.4)  thus  loses  6He  (=  6  X  4  =  24)  during  the  disin- 
tegration. There  are  thus  three  chlorides  of  lead  with  identical 
properties,  but  different  compositions,  namely  the  common  one 
207.2  :  2  X  35.46,  that  from  radium  206  :  2  X  35.46,  and  that 
from  thorium  208.4  :  2  X  35.46. 

Actinium  and  polonium  are  also  radioactive  elements,  which 
have  not  yet  been  fully  investigated.  The  former  appears  to  be 
formed  by  a  second,  parallel,  disintegration  of  Ui,  and  the  latter 
in  a  similar  way  from  Ra-E.  Compounds  of  potassium  and 
rubidium  show  traces  of  radioactivity. 

Significance  of  Radioactivity.  —  The  Brownian  movement 
(p.  416)  has  revealed  to  us  bodies  intermediate  between  ordinary 
particles  and  single  molecules,  and  has  enabled  us  to  estimate 
the  actual  weight  of  molecules.  Radioactivity  enables  us  to 
count  charged  molecules  of  helium  as  they  enter  the  electroscope 
or  produce  flashes  of  light  on  zinc  sulphide,  and  the  fog-tracks 
permit  us  to  follow  their  movements..  There  is  thus  now  no 
question  that  molecules  and  atoms  are  real.  Furthermore,  we 
infer  that  all  kinds  of  atoms  are  composed  of  a  positive  nucleus 
(p.  304)  surrounded  by  electrons,  although  only  the  atoms  of 


THE   RADIOACTIVE   ELEMENTS  615 

radioactive  elements  are  unstable.  The  diameter  of  the  positive 
nucleus  of  a  hydrogen  atom  is  calculated  to  be  about  T^^  of 
that  of  an  electron.  Rutherford  has  confirmed  this  by  actual 
measurement.  The  atom  is  thus  no  longer  regarded  as  being 
solid  and  continuous  in  structure.  It  is  mainly  a  vacuum,  con- 
taining a  few  relatively  very  minute  bodies  possessing  weight.  The 
fact  that  a-particles  are  thus  able  to  plough  their  way  through 
molecules  of  oxygen  and  nitrogen,  being  diverted  from  a  straight 
path  only  when  they  happen  to  pass  very  close  to  the  positive 
nucleus  (which,  of  course,  repels  the  positive  a-particles),  is  no 
longer  mysterious. 

Another  interesting  conclusion  has  been  reached  from  the  ob- 
servation that  niton  is  found  in  the  soil  and  in  many  natural 
waters.  Calculation  shows  that  the  heat  given  off  by  the  disin- 
tegration of  the  amounts  of  radioactive  matter  known  to  exist 
in  the  crust  of  the  earth  is  alone  sufficient  to  account  for  the 
maintenance  of  the  temperature  of  the  planet.  A  globe  of  the 
size  and  material  of  the  earth,  possessing  originally  only  heat 
energy,  and  cooling  from  a  white  hot  condition  to  the  temperature 
of  interstellar  space,  would  have  passed  through  the  stage  ^  of 
habitable  temperatures  in  a  much  shorter  time  than  that  which 
a  study  of  the  geological  deposits  (and  the  fossils  they  contain) 
show  to  have  been  actually  available.  The  discovery  of  the 
enormous,  but  gradually  released  disintegration  energy  of  the 
radioactive  elements  enables  us  now  to  explain  the  prolonged 
period  during  which  life  has  existed  on  the  earth. 

Exercises.  —  1.  Construct  equations,  showing  the  interactions  of: 
(a)  chromic  oxide  and  aluminium,  (6)  strontium  nitrate  and  potas- 
sium dichromate  in  solution,  (c)  potassium  hydroxide  and  chromic 
hydroxide,  and  the  reversal  on  boiling,  (d)  chlorine  and  potassium 
chromite  in  excess  of  alkali  (what  is  the  actual  oxidizing  agent?) . 

2.  What  volume  of  oxygen  at  0°  and  760  mm.,  (a)    is  obtain- 
able from  one  formula-weight  of  potassium  dichromate  (par.  4, 
p.  599),  (6)  is  required  to  oxidize  one  formula-weight  of  chromous 

chloride? 

3.  To  what  classes  of  actions  should  you  assign  the  three  met 

of  making  chromic  oxide  (p.  603)? 


616  COLLEGE    CHEMISTRY 

4.  Make  equations  for  all  the  reactions  involved  in  the  prepa- 
ration of  sodium-diuranate  from  pitchblende. 

5.  How  many  candle  power  will  be  obtained  from  50-watt 
carbon  and  tungsten  filament  lamps,  respectively? 

6.  Point  out  the  resemblance,  and  the  differences  between  the 
reactions  of,  (a)  gold  with  aqua  regia,  (b)  calcium  oxalate  with 
hydrochloric  acid,  (c)  barium  chromate  with  nitric  acid  (p.  601). 


CHAPTER  XLIV 
MANGANESE 

The  Chemical  Relations  of  the  Element.  —  Manganese 
stands,  at  present,  alone  on  the  left  side  of  the  eighth  column  of  the 
periodic  table.  The  right  side  is  occupied  by  the  halogens.  It  is 
never  univalent,  as  are  the  halogens,  but  its  heptoxide  Mn207  and 
the  corresponding  acid,  permanganic  acid  HMnO4,  are  in  many 
ways  closely  related  to  the  heptoxide  of  chlorine  and  perchloric 
acid  HC104.  Of  the  lower  oxides  of  manganese,  MnO  is  basic, 
and  Mn2Os  feebly  basic.  MnO2  is  feebly  acidic,  MnO3  more 
strongly  so,  and  permanganic  acid  (from  Mn20?)  is  a  very  active 
acid.  Contrary  to  the  habit  of  feebly  acidic  and  feebly  basic 
oxides,  such  as  those  of  zinc,  aluminium,  and  tin,  the  basic  oxides 
of  manganese  are  not  at  all  acidic,  and  the  acidic  oxides,  with  the 
exception  of  MnO2,  are  not  also  basic.  There  are  thus  the  five 
following,  rather  well-defined  sets  of  compounds,  showing  five 
different  valences  of  the  element.  Of  these  the  first,  fourth,  and 
fifth  are  the  most  stable  and  the  most  important. 

1.  Manganous    compounds,     MnO,     Mn(OH)2,     MnS04,     etc. 
These  compounds  resemble  those  of  the  magnesium  family  (and 
those  of  Fe++).     The  salts  of  weak  acids,  such  as  the  carbonate  and 
sulphide,  are  easily  made,  and  there  is  little  hydrolysis  of  the 
halides.     The  salts  are  pale-pink  in  color. 

2.  Manganic  compounds,  Mn2O3,  Mn(OH)3,  Mn2(S04)3,  [MnCl3]. 
The  salts  resemble  the  chromic  and  aluminium  salts  in  behavior, 
but  are  even  less  stable  than  those  of  quadrivalent  lead.     They  are 
completely  hydrolyzed  by  little  water.     The  salts  are  violet  in 
color. 

3.  Manganites,  Mn02,  H2Mn03,  CaMn03.     The  alkali  manga- 
nites  are  strongly  hydrolyzed,  like  the  plumbates  and  the  stannates. 

4.  Manganates,  Mn03,  H2Mn04,  K2Mn04.     The  salts  resemble 
the  sulphates  and  chromates,  but  are  much  more  easily  hydrolyzed. 
The  free  acid  resembles  chloric  acid  (p.  314)  in  that,  when  it  de- 

617 


618  COLLEGE    CHEMISTRY 

composes,  it  yields  a  higher  acid  (HMn04)  and  a  lower  oxide 
(Mn02).  The  salts  are  green  in  color. 

5.  Permanganates,  Mn2O7,  HMn04  (hydrated),  KMn04.  The 
salts  resemble  the  perchlorates,  and  are  not  hydrolyzed  by  water. 
They  are  reddish-purple  in  color. 

It  will  be  seen  that  the  element  manganese  changes  its  character 
totally  with  change  in  valence,  and  in  each  form  of  combination 
resembles  some  set  of  elements  of  valence  identical  with  that  which 
it  has  itself  assumed.  Since  the  valence  represents  the  number  of 
electrons  gained  or  lost  by  each  atom  (p.  322),  it  is  thus  evident 
that  the  chemical  properties  of  an  element  depend  more  upon 
the  electrical  constitution  of  its  atom  than  upon  the  atomic  weight. 
The  latter  is  a  secondary  property,  dependent  on  the  former 
(cf.  p.  304). 

Occurrence:  the  Metal.  —  The  chief  ore  is  the  dioxide,  pyro- 
lusite  MnO2,  which  always  contains  compounds  of  iron.  Other 
manganese  minerals  are:  braunite  Mn2O3;  the  hydrated  form, 
manganite  MnO(OH);  hausmannite  Mn3C>4;  and  manganese  spar 
MnC03.  The  metal  is  most  easily  made  by  reducing  one  of  the 
oxides  with  aluminium  by  Goldschmidt's  method. 

The  metal  manganese  (m.-p.  1260)  has  a  grayish  luster  faintly 
tinged  with  red.  It  is  oxidized  superficially  by  air,  and  easily  dis- 
places hydrogen  from  dilute  acids,  giving  manganous  salts.  Its 
alloys  with  iron,  such  as  spiegel  iron  (5-15  per  cent  Mn)  and  ferro- 
manganese  (70-80  per  cent  Mn),  are  made  by  using  manganese  ores 
with  the  charge  in  the  blast  furnace,  and  are  added  to  the  iron  in 
making  special  steels.  Manganese  steel  (7-20  per  cent  Mn)  is 
exceedingly  hard,  even  when  cooled  slowly.  It  is  used  for  the 
jaws  of  rock  crushing  machinery  and  for  burglar-proof  safes. 
Wires  made  of  an  alloy  called  manganin  (Cu  84  per  cent,  Ni  4  per 
cent,  Mn  12  per  cent),  invented  by  Weston,  is  used  in  instruments 
for  making  electrical  measurements,  because  its  resistance  does 
not  alter  with  moderate  changes  in  temperature. 

Oxides.  —  Manganous  oxide  MnO  is  a  green  powder,  made  by 
reducing  any  of  the  other  oxides  with  hydrogen.  Hausmannite 
Mn3O4  is  dull  red.  An  oxide  having  this  composition  is  formed 
when  any  of  the  other  oxides  is  heated  in  air,  oxidation  or  reduction, 


MANGANESE  619 

as  the  case  may  be,  taking  place  (cf.  p.  575).  Manganic  oxide 
Mn203  is  brownish-black,  and  is  formed  by  heating  any  of  the 
oxides  in  oxygen. 

Manganese  dioxide  Mn02  is  black,  and  is  most  easily  prepared 
in  pure  condition  by  gentle  ignition  of  manganous  nitrate.  The 
hydrated  forms  of  the  oxide  are  produced  by  precipitation,  as  by 
adding  a  hypochlorite  or  hypobromite  to  manganous  hydroxide 
suspended  in  water.  Manganese  dioxide  is  not  a  peroxide  in  the 
restricted  sense  (cf.  p.  223).  That  is  to  say,  it  does  not  contain  the 
radical  02n  and,  therefore,  does  not  give  hydrogen  peroxide.  Its 
reaction  formula  is  Mn(O)2  not  Mn(02)  and  in  double  decomposi- 
tions it  yields  only  water  H2(O).  It  is  used  for  manufacturing 
chlorine,  although  electrolytic  processes  are  now  driving  it  out  of 
this  field.  In  glass-making  (q.v.),  it  is  employed  to  oxidize  the 
green  ferrous  silicate,  derived  from  impurities  in  the  sand,  to  the 
pale-yellow  ferric  compound.  The  amethyst  color  of  the  manganic 
silicate  which  is  formed  tends  also  to  neutralize  this  yellow.  It  is 
mixed  with  black  paints  as  a  " dryer"  (oxidizing  agent). 

Manganese  trioxide  Mn03  is  a  red,  unstable  powder.  Manganese 
heptoxide  Mn2O7  is  a  brownish-green,  volatile  oil  (see  below). 

When  any  of  these  oxides  is  heated  with  an  acid,  a  manganous 
salt  is  obtained.  Salts  of  this  class  are,  in  fact,  the  only  stable  sub- 
stances in  which  manganese  is  combined  with  an  acid  radical.  In 
this  action  the  oxides  containing  more  oxygen  than  does  MnO  give 
off  oxygen,  or  oxidize  the  acid  (cf.  p.  157).  When  the  oxides  are 
heated  with  bases,  in  the  presence  of  air,  manganates  are  always 
formed.  In  this  case,  with  oxides  containing  a  smaller  proportion 
of  oxygen  than  MnOa,  oxygen  is  taken  from  the  air. 

Manganous  Compounds.  —  The  manganous  salts  are  formed 
by  the  action  of  acids  upon  the  carbonate  or  any  of  the  oxides. 
Thus  the  chloride  MnCl2,4H2O  is  obtained  in  pale-pink  crystals 
from  a  solution  made  by  treating  the  dioxide  with  hydrochloric  acid 
and  driving  off  the  chlorine  liberated  by  oxidation  (p.  158).  The 
hydroxide  Mn(OH)2  is  formed  as  a  white  precipitate  when  a  soluble 
base  is  added  to  a  solution  of  a  manganous  salt.  This  body  passes 
into  solution  when  ammonium  salts  are  added,  and  cannot  be 
precipitated  in  their  presence  on  account  of  the  formation  of 
molecular  ammonium  hydroxide  and  the  suppression  of  hydroxide- 


620  COLLEGE    CHEMISTRY 

ion  (cf.  magnesium  hydroxide,  p.  525).  The  hydroxide  quickly 
darkens  when  exposed  to  the  air  and  passes  over  into  hydrated 
manganic  oxide  MnO(OH). 

Manganous  sulphate  gives  pink  crystals  of  a  hydrate.  Below  6° 
the  solution  deposits  MnSO4,7H20,  which  is  a  vitriol  (p.  529). 
Between  7°  and  20°  the  product  is  MnSO4,5H20,  asymmetric  and 
resembling  CuSO4,5H20.  Above  25°  monosymmetric  prisms  of 
MnS04,4H2O  are  obtained.  These  hydrates  have  different  aqueous 
tensions  and  may  be  formed  from  one  another  by  lowering  or  raising 
the  pressure  of  water  vapor  around  the  substance  (p.  96). 

Manganous  carbonate  MnCO3  is  a  white  powder  formed  by  pre- 
cipitation. The  sulphide  MnS  is  obtained  as  a  green,  crystalline 
powder  by  leading  hydrogen  sulphide  over  any  of  the  oxides.  A 
flesh-colored,  amorphous  manganous  sulphide  MnS  (often  some- 
what hydrated)  is  more  familiar  and  is  precipitated  by  ammonium 
sulphide  from  manganous  salts.  It  interacts  with  mineral  acids 
and  even  with  acetic  acid,  so  that  it  cannot  be  precipitated  by  the 
action  of  hydrogen  sulphide  on  salts  (cf.  p.  530).  When  rubbed  in 
a  mortar  it  becomes  crystalline,  and  is  then  green. 

The  manganous  salts  of  weak  acids,  such  as  the  carbonate  and 
sulphide,  darken  when  exposed  to  air  and  are  oxidized,  with  forma- 
tion of  hydrated  manganic  oxide.  As  we  have  seen,  manganous 
hydroxide  is  similarly  oxidized  and  these  salts  are  precisely  the  ones 
which  should  furnish  the  hydroxide  by  hydrolysis.  While  there  is 
a  general  resemblance  between  the  manganous  salts  and  the  stan- 
nous,  chromous,  and  ferrous  salts,  the  manganous  salts  of  active 
acids  are  not  oxidized  by  the  air  as  are  the  corresponding  salts  of 
the  other  three  metals. 

Manganic  Compounds.  —  The  base  of  this  set  of  compounds, 
manganic  hydroxide  Mn(OH)3,  is  slowly  deposited  by  the  action  of 
the  air  on  an  ammoniacal  solution  of  a  manganous  salt  in  salts  of 
ammonium.  Manganic  chloride  MnCls  is  present  in  the  liquid  ob- 
tained by  the  action  of  hydrochloric  acid  upon  manganese  dioxide 
(cf.  p.  158),  but  loses  chlorine  very  readily. 

Manganites.  —  Although  manganese  dioxide  interacts  when 
fused  with  potassium  hydroxide,  simple  salts  derived  from 
H2MnO3  (=  H20,MnO2)  or  H4MnO4  ( =  2H20,Mn02)  are  not 


MANGANESE  621 

formed.  The  products  are  complex,  as  K2Mn5On.  Some  less 
complex  manganites  are  formed  by  mixing  manganous  chloride 
solution  with  slaked  lime,  and  blowing  air  through  the  mass  of 
calcium  and  manganous  hydroxides  which  is  thus  obtained.  Man- 
ganites of  calcium,  such  as  CaMnO3  ( =  CaO,MnO2)  and  CaMn205 
(=  CaO,2Mn02)  are  thus  formed: 

Ca(OH)2  +  2Mn(OH)2  -f  O2  ->  CaMn205  +  3H2O. 

Manganates.  —  When  one  of  the  oxides  of  manganese  is  fused 
with  potassium  carbonate  and  potassium  nitrate,  a  green  mass  is 
obtained.  The  green  aqueous  extract  deposits  potassium  man- 
ganate  K2MnO4  in  rhombic  crystals,  which  are  of  the  same  form  as 
those  of  potassium  sulphate,  and  are  almost  black: 

K2C03  +  MnO2  +  O  ->  K2MnO4  +  CO2. 

The  acid  H2Mn04  is  itself  unknown.  The  potassium  salt  remains 
unchanged  in  solution  only  in  presence  of  free  alkali.  When  the 
concentration  of  the  hydroxide-ion  is  reduced  by  dilution,  or,  better 
still,  when  a  weak  acid  such  as  carbonic  acid  or  acetic  acid  is  used 
to  neutralize  it,  the  salt  is  decomposed  according  to  the  following 
equation : 

3K2Mn04  +  2H20  ->4KOH  +  2KMn04  +  Mn02. 

That  is,  a  precipitate  of  manganese  dioxide  and  a  solution  of 
potassium  permanganate  are  obtained.  To  make  the  equation 
(pp.  322-324),  we  note  that  in  K2MnO4  we  have  2K+  and  4O=  and 
therefore  Mn+tt  to  secure  electrical  neutrality.  The  latter  becomes 
Mnfl^  and  Mntt  Arithmetically  3MnW  will  give  2MnWf 
and  iMntt  Hence,  3K2MnO4  are  required,  and  2KMnO4  and 
!MnO2  produced.  In  terms  of  the  ions  the  equation  is  simpler: 

3MnO4=  +  2H+  -»  20ET  +  2Mn04~  +  MnO2. 

Permanganates.  —  Potassium  permanganate  KMn04  is  made 
by  decomposition  of  the  manganate  as  shown  above,  and  is  ob- 
tained, in  purple  crystals  with  a  greenish  luster,  by  evaporation  of 
the  solution.  To  avoid  the  loss  of  manganese  thrown  down  as 
dioxide,  the  action  is  carried  out  commercially  by  passing  ozone 


622  COLLEGE    CHEMISTRY 

through  the  solution  of  the  manganate:  2K2MnO4  -f  O3  -j- 
H2O  —  >  2KMn04  +  02  +  2KOH.  Sodium  permanganate  NaMn04 
is  made  in  a  similar  manner.  It  is  not  obtainable  in  solid  form,  but 
its  solution  is  known  as  "Condy's  disinfecting  fluid."  This  liquid 
owes  its  properties  to  the  oxidizing  power  of  the  salt.  Perman- 
ganic acid  is  a  very  active  acid,  that  is,  it  is  highly  ionized  in 
aqueous  solution.  A  solid  hydrate  of  the  acid  may  be  secured  in 
reddish-brown  crystals  by  adding  sulphuric  acid  to  a  solution  of 
barium  permanganate  and  allowing  the  filtrate  to  evaporate  : 
Ba(Mn04)2  +  H2S04  +  zH20  <=±BaS04  j  +  2HMn04,xH20. 

This  hydrate  decomposes,  on  being  warmed  to  32°,  and  yields 
oxygen  and  manganese  dioxide.  When  a  very  little  dry,  powdered 
potassium  permanganate  is  moistened  with  concentrated  sulphuric 
acid,  brownish-green,  oily  drops  of  permanganic  anhydride  (man- 
ganese heptoxide)  Mn2O?  are  formed.  This  compound  is  volatile, 
giving  a  violet  vapor,  and  is  apt  to  decompose  explosively  into 
oxygen  and  manganese  dioxide.  Its  oxidizing  power  is  such  that 
combustibles  like  paper,  ether,  and  illuminating  gas  are  set  on  fire 
by  contact  with  it. 

Potassium  Permanganate  as  an  Oxidising  Agent.  —  The 

actions  are  different  according  as  the  substance  is  employed  (1)  in 
acid,  or  (2)  in  neutral  solution. 

1.  In  presence  of  an  acid,  and  an  oxidizable  body,  a  manganous 
salt  is  always  formed.  The  schematic  equation,  Mn2C>7—  »2MnO-f 
5O,  shows  that  every  two  molecules  of  the  permanganate  yield  50 
for  oxidizing  purposes.  Thus,  when  sulphuric  acid  is  added  to 
potassium  permanganate  solution,  and  sulphur  dioxide  is  led 
through  the  mixture,  we  have: 

2KMn04  +  3H2S04  -»  K2S04+2MnS04+3H2O(+50)   (1) 
5H2SO3-»5H2SO4  (2) 


_ 

2KMnO4+3H2SO4+  5H2SO3  -»  K2SO4+  2MnSO4  +  3H2O  +  5H2SO4 

In  this  case,  since  sulphuric  acid  is  a  product,  the  preliminary  addi- 
tion of  the  acid  was  superfluous.  In  other  cases,  the  partial  equa- 
tion (1),  showing  the  available  50,  remains  the  same,  while  the 
other  partial  equation  varies  with  the  substance  being  oxidized. 
Thus,  with  hydrogen  sulphide  as  reducing  agent,  we  have: 

(0)+H2S-+H20  +  S  X5  (20 


MANGANESE  623 

and  with  ferrous  sulphate,  we  get  ferric  sulphate: 

2FeS04  +  H2S04(+  O)  f+  Fe2(S04)3  +  H2O       X  5         (2") 

As -before  (2')  and  (2")  must  be  multiplied  throughout  by  five, 
before  summation  is  made  (see  also  p.  225) . ' 

The  quantity  of  a  ferrous  salt,  or  of  hydrogen  peroxide  (p.  225) 
in  a  sample  of  a  solution  may  be  measured  by  titrating  (p.  257)  the 
solution  with  a  standard  solution  of  potassium  permanganate  until 
the  color  ceases  to  be  destroyed,  and  then  noting  the  volume  used. 
For  iron,  the  standard  solution  may  be  prepared  so  that  1  cc.  will 
oxidize  0.01  g.  of  Fe++. 

2.  When  dry  potassium  permanganate  is  heated,  it  decomposes 
as  follows: 

2KMn04  -*  K2Mn04  +  MnO2  +  O2. 

The  neutral  solution  oxidizes  substances  which  are  reducing  agents. 
The  fingers  are  stained  brown  by  an  aqueous  solution,  receiving  a 
deposit  of  manganese  dioxide,  in  consequence  of  the  reducing 
power  of  the  unstable  organic  substances  in  the  skin.  The  de- 
struction of  minute  organisms  by  Condy's  fluid  results  from  a 
similar  action.  When  the  powdered  salt  is  moistened  with  glycer- 
ine, the  mass  presently  bursts  into  flame. 

Analytical  Reactions  of  Manganese  Compounds.  —  The 

ions  commonly  encountered  are  manganous-ion  Mn++,  which  is 
very  pale-pink  in  color,  permanganate-ion  MnO4~,  which  is  purple, 
and  manganate-ion  Mn04=,  which  is  green.  The  manganous 
compounds  give  with  ammonium  sulphide  the  flesh-colored  sul- 
phide which  is  soluble  in  acids.  Bases  give  the  white  hydroxide, 
which  darkens  by  oxidation,  and  is  soluble  in  salts  of  ammonium. 
All  compounds  of  manganese  confer  upon  the  borax  bead  an 
amethyst  color  (manganic  borate),  which,  in  the  reducing  flame, 
disappears  (manganous  borate).  A  bead  of  sodium  carbonate 
and  niter  becomes  green  on  account  of  the  formation  of  the 
manganate. 

Exercises.  —  1.  What  do  we  mean  by  saying  that,  (a)  chromous 
chloride  is  stable  (p.  93),  but  easily  oxidized  by  the  air,  (6)  per- 


624  COLLEGE    CHEMISTRY 

manganic  acid  is  an  active  oxidizing  agent  in  presence  of  an  acid 
(p.  622). 

2.  Formulate  the  oxidations  of  hydrogen  sulphide,  of  ferrous 
sulphate,  of  oxalic  acid  (to  carbon  dioxide),  and  of  nitrous  acid  (to 
nitric  acid)  by  potassium  permanganate  in  acid  solution.  In  doing 
so,  employ  the  several  methods  suggested  on  pp.  322-326. 


CHAPTER  XLV 
IRON,   COBALT,   NICKEL 

THE  elements  iron  (Fe,  at.  wt.  55.84),  cobalt  (Co,  at.  wt.  59),  and 
nickel  (Ni,  at.  wt.  58.7)  are  not  corresponding  members  of  succes- 
sive periods,  like  the  families  hitherto  considered.  They  are 
neighboring  members  of  the  first  long  period,  lying  between  its  first 
and  second  octaves. 

IKON  Fe 

Chemical  Relations  of  the  Element.  —  The  oxides  and 
hydroxides  FeO  and  Fe(OH)2,  Fe203  and  Fe(OH)3  are  basic,  the 
former  more  strongly  so  than  the  latter.  The  ferrous  salts,  de- 
rived from  Fe(OH)2,  resemble  those  of  the  magnesium  group  and 
those  of  Cr++  and  Mn++,  and  are  little  hydrolyzed.  The  ferric 
salts,  derived  from  Fe(OH)3,  resemble  those  of  Cr+++  and  A1+++ 
and  are  hydrolyzed  to  a  considerable  extent.  Iron  gives  also  a 
few  ferrates  K2FeO4,  CaFeO4,  etc.,  derived  from  an  acid  H2Fe04 
which,  like  manganic  acid  H2Mn04  (p.  621),  is  too  unstable  to  be 
isolated.  Complex  anions  containing  iron,  such  as  the  anion  of 
K4.Fe(CN)6,  are  familiar,  but  complex  cations  containing  ammonia 
are  unknown. 

Occurrence.  —  Free  iron  is  found  in  minute  particles  in  some 
basalts,  and  many  meteorites  are  composed  of  it.  Meteoric  iron 
can  be  distinguished  from  specimens  of  terrestrial  origin  by  the 
fact  that  it  contains  3-8  per  cent  of  nickel.  The  chief  ores  of  iron 
are  the  oxides,  haematite  Fe203  and  magnetite  Fe3O4,  and  the  car- 
bonate FeCO3,  siderite.  The  first  is  reddish  and  radiated  in 
structure;  but  black,  shining,  rhombohedral  crystals,  known  as 
specularite,  are  also  found.  Hydrated  forms,  like  brown  iron  ore 
2Fe2O3,3H20,  are  also  common.  Siderite  is  pale-brown  in  color 
and  rhombohedral,  like  calcite.  When  mixed  with  clay  it  forms 
iron-stone,  from  which  most  of  the  iron  in  Great  Britain,  but  less 
than  one  per  cent  of  that  in  the  United  States  is  obtained.  Pyrite 

625 


626 


COLLEGE    CHEMISTRY 


consists  of  golden-yellow,  shining  cubes  or  pentagonal  dodec- 
ahedra.  It  is  used,  on  account  of  its  sulphur,  in  the  manufacture 
of  sulphuric  acid,  but,  from  the  oxidized  residue,  iron  of  sufficient 
purity  is  obtained  with  difficulty.  Compounds  of  iron  are  con- 
tained in  chlorophyll  and  in  the  blood  (haemoglobin),  and  doubtless 
play  an  important  part  in  connection  with  the  vital  functions  of 
these  substances.  Ammonium  sulphide  blackens  the  skin,  form- 
ing ferrous  sulphide  by  interaction  with  organic  compounds  of  iron 
present  in  the  tissues. 

Pure  Iron.  —  Pure  iron  is  obtained  by  reducing  pure  ferrous 
oxalate  in  a  stream  of  hydrogen  at  a  high  temperature.  It  is  also 
made  by  electrolysis  of  ferrous  sulphate  solution  at  100°  between 
iron  electrodes.  It  is  silver-white  and  melts  at  1510°.  The  purest 
iron  does  not  rust  in  pure  cold  water,  but  the  impurities  in  ordinary 
iron  act  as  contact  agents  and  rusting  proceeds. 

Metallurgy.  —  The  ores  of  iron  are  first  roasted  in  order  to 
decompose  carbonates  and  oxidize  sulphides,  if  these  salts  are 
present.  Coke  is  then  used  to  reduce  the  oxides. 
Coal  is  unsuitable  because  so  much  heat  is  wasted 
in  driving  out  the  volatile  matter  and  moisture, 
which  are  absent  from  coke.  Ores  containing  lime 
or  magnesia  are  mixed  with  an  acid  flux,  such  as 
sand  or  clay-slate,  in  order  that  a  fusible  slag  may 
be  formed.  Conversely,  ores  containing  silica  and 
clay  are  mixed  with  limestone.  With  proper 
adjustment  of  the  ingredients  the  process  can  be 
carried  on  continuously  in  a  blast  furnace  (Fig. 
136),  an  iron  structure  40  to  100  feet  high,  lined 
with  firebrick.  The  solid  materials  thrown  in  at 
the  top  are  converted,  as  they  slowly  descend, 
completely  into  gases  which  escape  and  liquids 
(iron  and  slag)  which  are  tapped  off  at  the  bot- 
tom. Heated  air  is  blown  in  at  the  bottom  through 
tuyeres,  and  the  top  is  closed  by  a  cone  which  descends  for  a  moment 
when  an  addition  is  made  to  the  charge.  The  gases,  which  contain 
much  carbon  monoxide,  are  led  off  and  used  to  heat  the  blast  or  to 
drive  gas-engines. 


Fia.  136. 


IRON  627 

The  main  action  takes  place  between  the  carbon  monoxide, 
present  in  consequence  of  the  excess  of  carbon,  and  the  oxide  of 
iron: 

Fe3O4  +  4CO  <=»  3Fe  +  4CO2. 

Since  the  action  is  a  reversible  one,  a  large  excess  of  carbon  mon- 
oxide is  required.  At  650°,  equilibrium  is  reached  with  CO  :  C02  :: 
1  vol.  :  lj  vols.,  and  in  practice  the  proportion  of  carbon  monoxide 
used  is  from  twice  to  fifteen  times  as  great.  Almost  5  tons  of  air, 
heated  in  advance  to  800°,  are  required  for  each  ton  of  iron 
produced.  The  moisture  in  this  air  acts  upon  the  coke,  giving 
water-gas  (p.  386).  This  action  uses  up  fuel,  and  also  lowers  the 
temperature  at  the  point  where  it  should  be  highest.  In  the  most 
modern  furnaces,  therefore,  the  air  is  dried  (Gayley  dry-blast 
process),  with  a  saving  in  coke  equivalent  to  SI. 00  (4/-)  per  ton 
of  iron  obtained.  This  illustrates  the  commercial  value  of  even 
a  single  improvement  in  a  chemical  operation.  If  the  Gayley 
process  were  used  with  every  blast  furnace,  an  immense  sum  would 
be  saved,  for  in  the  United  States  alone  30  million  tons  of  iron  are 
annually  produced  (1913).  This  is  considerably  over  40  per  cent 
of  the  world's  production,  20  per  cent  being  supplied  by  Germany 
and  15  per  cent  by  Great  Britain. 

In  the  upper  part  of  the  furnace,  the  heat  (400°)  loosens  the 
texture  of  the  ore.  Further  down,  the  temperature  is  higher 
(500-900°),  and  the  carbon  monoxide  reduces  the  oxide  of  iron  to 
particles  of  soft  iron.  A  temperature  high  enough  to  melt  pure 
iron  is  barely  reached  anywhere  in  the  furnace,  but,  a  little  lower 
down,  by  solution  of  carbon  in  the  iron,  the  more  fusible  cast  iron 
(m.-p.  about  1200°)  is  formed  and  falls  in  drops  to  the  bottom. 
It  is  in  this  region  also  that  the  slag,  essentially  a  glass  (p.  493), 
is  produced.  If  the  flux  had  begun  sooner  to  interact  with  the 
unreduced  ore,  iron  would  have  been  lost  by  the  formation  of  the 
silicate.  The  iron  collects  below  the  slag,  and  the  latter  flows 
continuously  from  a  small  hole.  The  former  is  tapped  off  at 
intervals  of  six  hours  or  so  from  a  lower  opening.  As  a  rule,  the 
iron  never  cools  until  it  has  been  converted  into  rails  or  structural 
iron.  In  some  cases,  it  is  made  into  "pigs"  in  a  casting  machine. 

Cast  Iron  and  Wrought  Iron.  —  Pure  iron  is  not  manu- 
factured, and  indeed  would  be  too  soft  for  most  purposes.  Piano- 


628  COLLEGE    CHEMISTRY 

wire,  however,  is  about  99.7  per  cent  pure.  The  product  obtained 
from  the  blast  furnace  contains  92-94  per  cent  of  iron  along  with 
2.6-^.3  per  cent  of  carbon,  often  nearly  as  much  silicon,  varying 
proportions  of  manganese,  and  some  phosphorus  and  sulphur. 
The  last  four  ingredients  are  liberated  from  combination  with 
oxygen  by  the  carbon  in  the  hottest  part  of  the  furnace  and  com- 
bine or  alloy  themselves  with  the  iron.  Cast  iron  does  not  soften 
before  melting,  as  does  the  purer  wrought  iron  (m.-p.  1510°),  but 
melts  sharply  at  1150-1250°  according  to  the  amount  of  foreign 
material  it  contains.  When  suddenly  cooled  it  gives  chilled  cast 
iron  which  is  very  brittle  and  looks  homogeneous  to  the  eye,  all  the 
carbon  being  present  in  the  form  of  carbide  of  iron  Fe3C  (cementite) 
in  solid  solution  in  the  metal.  This  solid  solution  is  exceedingly 
hard,  but  very  brittle.  By  slower  cooling,  time  is  permitted  for 
the  separation  of  part  of  the  carbon  as  graphite,  which  appears  in 
tiny  black  scales,  and  gray  cast  iron  results.  This  mixture  is  much 
softer,  on  account  of  the  amount  of  free,  relatively  pure  iron  which 
it  contains. 

Cast  iron  is  used  in  making  cooking  ranges,  stoves,  pipes,  and 
radiators.  It  expands  in  solidifying,  and  so  fills  every  detail  of  the 
mold. 

Wrought  iron,  invented  by  Henry  Cort  (1784),  is  made  by  heat- 
ing the  broken  pigs  of  cast  iron  upon  a  layer  of  material  containing 
oxide  of  iron  and  hammer-slag  (basic  silicate  of  iron)  spread  on  the 
bed  of  a  reverberatory  furnace  (Fig.  116,  p.  460).  The  carbon, 
silicon,  and  phosphorus  combine  with  the  oxygen  of  the  oxide,  and 
the  last  two  pass  into  the  slag.  The  sulphur  is  found  in  the  slag 
as  ferrous  sulphide.  On  account  of  the  effervescence  due  to  the 
escape  of  carbon  monoxide,  the  process  is  called  " pig-boiling." 
The  iron  is  stirred  with  iron  rods  (" puddled")  and  stiffens  as  it 
becomes  purer,  until  finally  it  can  be  withdrawn  in  balls  (" blooms") 
and  partially  freed  from  slag  by  rolling.  The  resulting  bars  are 
repeatedly  cut,  piled  in  a  bundle,  reheated,  and  rolled.  The  iron 
now  softens  sufficiently  for  welding  below  1000°  and  melts  at 
1505°.  Its  fibrous  structure  is  due  partly  to  the  films  of  slag  which 
have  not  been  completely  pressed  out  by  the  rolling.  On  account 
of  its  toughness,  wrought  iron  is  used  for  anchors,  chains,  and  bolts, 
and  for  drawing  into  wire.  On  account  of  its  relative  purity 
(99.8-99.9  per  cent),  it  is  less  fusible  than  cast  iron  and  is  used  for 


IRON  629 

fire  bars.  The  above  operations  are  now  largely  performed  by 
machinery,  but  have  been  largely  displaced  by  the  Bessemer  and 
open  hearth  processes  in  which  iron  of  equal  purity  can  be  obtained. 

Properties  of  Steel.  —  This  is  a  variety  of  iron  almost  free 
from  phosphorus,  sulphur,  and  silicon.  Tool-steel  contains  0.9- 
1.5  per  cent  of  carbon,  structural  steel  only  0.2-0.6  per  cent,  and 
mild  steel  0.2  per  cent  or  even  less.  Steel  combines  the  properties 
of  cast  and  of  wrought  iron,  being  hard  and  elastic,  and  at  the  same 
time  available  for  forging  and  welding  when  the  proportion  of 
carbon  is  not  too  high.  Steel  can  be  tempered  (see  below).  It 
has  also  a  greater  tensile  strength  *  than  has  wrought  iron,  and  it 
can  be  permanently  magnetized. 

Bessemer  Process.  —  Steel  is  made  largely  by  the  Bessemer 
process  (Kelly  1852,  Bessemer  1855).     The  molten  cast  iron  is 
poured  into  a  converter  (Fig.  137)  and  a  blast  of  air  (a)  is  blown 
through  it.     The  oxidation  of  the 
manganese,    carbon,    silicon,    and 
a  little  of  the  iron  gives  out  suffi- 
cient heat  to  raise  the  temperature 
of  the  mass  above  the  melting- 
point  of  wrought   iron.     The  re- 
quired  proportion    of    carbon    is 
then  introduced  by  adding  pure 
cast  iron,  spiegel  iron,  or  coke,  and  FIQ  137 

the  contents,  first  the  slag,  and 

then  the  molten  steel,  are  finally  poured  out  by  turning  the  con- 
verter. When  the  cast  iron  contains  much  phosphorus,  the  oxide 
of  this  element  is  reduced  again  by  the  iron  as  fast  as  it  is  formed 
by  the  blast.  In  such  cases  a  basic  lining  containing  lime  and 
magnesia  takes  the  place  of  the  sand  and  clay  lining  of  the  ordinary 
Bessemer  converter,  and  a  slag  containing  a  basic  phosphate  of 
calcium  is  produced.  This  modification  constitutes  what  is 
known  as  the  basic  or  Thomas-Gilchrist  process.  The  slag 
(" Thomas-slag")  when  pulverized  forms  a  valuable  fertilizer 

*  Tensile  strength  or  tenacity  is  measured  by  the  weight  (in  kilos)  required  to 
break  a  wire  of  the  metal  1  sq.  mm.  in  section.  Lead  2.6,  copper  51,  iron  71, 
steel  91. 


630 

(cf.  p.  488). 
preferred. 


COLLEGE    CHEMISTRY 


In  the  United  States,  the  basic  open-hearth  process  is 


Open-Hearth  (Siemens-Martin)  Process.  —  In  this  process 
the  cast  iron  is  melted  in  a  saucer-shaped  depression  (Fig.  138), 
which  is  lined  with  sand  in  the  acid  process  and  with  lime  and 
magnesia  in  the  basic  process.  Scraps  of  iron  plate  (for  dilution) 
and  haematite,  or  some  other  oxide  ore,  are  then  added  in  proper 
proportions.  The  materials  (50—75  tons  in  one  charge)  are 
heated  with  gas  fuel  for  8-10  hours.  To  secure  economically  the 


FIG.  138. 

high  temperature  required  to  keep  the  product  (almost  pure  iron) 
fused,  Siemens  devised  the  method  of  preheating  the  fuel  gas  and 
air  by  a  regenerative  device.  The  spent  air  and  gas  pass  down 
through  a  checkerwork  of  brick.  When  this  becomes  heated,  the 
valves  are  reversed,  the  gas  and  air  now  enter  through  the  heated 
brickwork  and,  after  meeting  and  burning  over  the  iron,  pass  out 
through  the  checkerwork  on  the  opposite  side,  raising  its  tempera- 
ture in  turn. 

The  changes  are  similar  to  those  in  the  Bessemer  process. 
During  casting,  some  aluminium  is  added  to  combine  with  oxygen 
(present  as  CO)  and  give  sounder  ingots.  Recently,  iron  con- 
taining 10-15  per  cent  of  titanium  has  been  added  instead.  The 


IRON       .  631 

titanium  combines  with  both  nitrogen  and  oxygen  and  the  com- 
pounds pass  into  the  slag,  just  as  does  aluminium  oxide.  Rails 
made  of  steel  purified  with  this  element  are  less  liable  to  breakage 
(the  commonest  cause  of  wrecks)  and  are  40  per  cent  more  durable, 
than  are  ordinary  open-hearth  rails. 

The  advantage  of  the  open-hearth  process  over  that  of  Bessemer 
is  that  it  is  not  hurried,  and  is  therefore  under  better  control.  The 
material  can  be  tested  by  sample  at  intervals  until  the  required 
composition  has  been  reached.  The  product  is  of  more  uniform 
quality.  When  fine  steel  is  required,  electric  heating  (e.g.,  in  the 
Heroult  furnace)  permits  even  more  deliberate  treatment. 

Bessemer  and  open-hearth  steel  is  used  for  heavy  and  light 
machinery  castings  and  for  shafts.  It  is  rolled  into  rails,  and  into 
bridge  and  structural  iron. 

Crucible  Steel.  —  For  special  purposes  steel  is  made  in  cru- 
cibles of  clay  (or  graphite  and  clay)  in  melts  of  60-100  pounds. 
"  Melting  bar,"  a  very  pure  open-hearth  steel,  is  melted  with 
charcoal  or  with  pure  pig  iron.  This  steel  is  employed  in  making 
razors  (1.5  per  cent  C),  tools  (1  per  cent  C),  dies  (0.75  per  cent  C), 
pens,  needles,  and  cutlery. 

Tempering.  —  The  carbon  in  steel  (and  cast  iron)  is  in  the 
form  of  carbon  or  of  carbide  of  iron  Fe3C  (6.6  per  cent  C),  dissolved 
in  the  iron.  When  white  hot  steel  (up  to  2  per  cent  C)  is  suddenly 
chilled,  there  is  no  time  for  any  changes  to  occur  during  the  cooling, 
and  a  solid  solution  is  obtained  which  is  very  hard  and  brittle. 
When,  however,  the  cooling  is  slow,  some  of  the  carbon  separates 
in  minute  crystals  of  cementite  FesC  until,  at  about  700°,  there 
remains  only  about  0.9  per  cent  carbon  in  solid  solution.  At  this 
temperature,  if  sufficient  time  is  allowed,  the  solid  solution  sepa- 
rates into  a  mixture  of  pure  iron  (87  per  cent)  which  is  soft  and 
carbide  of  iron  (13  per  cent)  which  is  hard.  Steel,  when  slowly 
cooled,  is  thus  a  mixture,  and  not  homogeneous.  If,  therefore, 
hard  chilled  steel  is  heated  once  more  for  the  purpose  of  tempering, 
the  extent  to  which  the  softer  material  is  formed  depends  upon  the 
temperature  reached  and  upon  the  rate  and  the  duration  of  the 
cooling  process.  By  varying  these,  the  degree  of  hardness  allowed 
to  remain  can  be  adjusted. 


632  COLLEGE    CHEMISTRY 

Steel  Alloys.  —  As  we  have  seen,  substances  such  as  aluminium, 
titanium,  and  ferrosilicon  are  added  to  iron  for  the  purpose  of 
purifying  it,  and  pass  in  combination  into  the  slag.  There  are, 
however,  regular  alloys  containing  the  foreign  metal  along  with 
the  iron.  Thus,  manganese  steel  (7-20  per  cent  Mn),  made  by 
adding  spiegel  iron  or  ferromanganese  (p.  618)  to  steel,  remains 
hard  even  when  cooled  slowly  and  is  used  for  the  jaws  of  rock- 
crushers  and  for  safes.  Chromium-vanadium  steel  (1  per  cent  Cr, 
0.15  per  cent  Va)  has  great  tensile  strength,  can  be  bent  double 
while  cold,  and  offers  great  resistance  to  changes  of  stress  and  to 
torsion.  It  is  used  for  frames  and  axles  of  automobiles  and  for 
connecting  rods.  Tungsten  steel  has  already  been  described  (p. 
606).  Nickel  steel  (2-4  per  cent  Ni)  resists  corrosion,  has  a  high 
limit  of  elasticity  and  great  hardness,  and  is  used  for  armor-plato, 
wire  cables,  and  propeller  shafts.  Invar  (36  per  cent  Ni)  is 
practically  non-expansive  when  heated  within  narrow  limits  and 
is  used  for  meter-scales  and  pendulum  rods. 

Chemical  Properties  of  Iron.  —  Although  the  purest  iron 
does  not  rust  in  cold  water  (p.  626),  ordinary  iron  rusts  in  moist 
air  or  under  water.  It  probably  rusts  in  water  free  from  carbon 
dioxide,  displacing  the  hydrogen-ion,  but  the  action  is  greatly 
hastened  by  the  presence  of  carbonic  acid.  Rust  is  a  brittle,  porous, 
loosely  adherent  coating  of  variable  composition,  consisting  mainly 
of  a  hydrated  ferric  oxide  3Fe2O3,H2O,  which  does  not  protect  the 
metal  below.  Oil  protects  iron  from  rusting  because,  although 
oxygen  is  more  soluble  in  most  oils  than  in  water,  and  so  reaches  the 
iron  freely,  water  is  not  soluble  in  oil  and  so  moisture  is  excluded. 

Iron  burns  in  oxygen  and  it  interacts  with  superheated  steam, 
in  both  cases  giving  Fe3O4.  A  superficial  layer  of  this  oxide  ad- 
heres firmly  and  protects  the  iron  from  the  action  of  the  air  (Barff's 
process  iron,  or  Russia  iron). 

Iron  displaces  hydrogen  easily  from  dilute  acids.  Steel  and  cast 
iron,  which  contain  iron,  its  carbide,  and  graphite,  give  with  cold 
dilute  acids  almost  pure  hydrogen,  and  the  carbide  and  graphite 
remain  unattacked.  More  concentrated  acids,  however,  particu- 
larly when  warm,  generate,  along  with  hydrogen,  hydrocarbons 
formed  by  interaction  with  the  carbide  (p.  441).  The  odor  of  the 
gas  is  due  to  compounds  of  sulphur  and  phosphorus. 


IRON  633 

Although  iron  acts  vigorously  on  dilute  or  concentrated  nitric 
acid,  it  is  indifferent  to  fuming  nitric  acid  (N02  in  solution,  p.  348). 
It  becomes  passive.  In  this  state,  it  no  longer  displaces  hydrogen 
from  dilute  acids.  If  dipped  in  cupric  sulphate  solution,  it  does 
not  receive  the  usual  red  coating  of  metallic  copper.  However,  if 
scratched  or  struck,  the  passive  condition  is  destroyed,  and  copper 
begins  to  be  deposited  at  the  point  touched  and  the  action  spreads 
quickly  over  the  whole  surface.  No  satisfactory  explanation  of 
this  phenomenon  has  been  obtained,  although  it  is  shown  also  by 
chromium,  cobalt,  and  other  metals. 

Ferrous  Compounds.  —  Ferrous  chloride  is  obtained  as  a  pale- 
blue  hydrate  FeCl2,4H2O  (turning  green  in  the  air)  by  interaction 
of  hydrochloric  acid  with  the  metal  or  the  carbonate.  The  an- 
hydrous salt  sublimes  in  colorless  crystals  when  hydrogen  chloride 
is  led  over  the  heated  metal.  In  solution  the  salt  is  oxidized  by 
the  air  to  a  basic  ferric  chloride: 

4Fe++  +  02  +  2H2O  ->  4Fe+++  +  4OET. 

In  presence  of  excess  of  the  acid,  normal  ferric  choride  is  formed. 
With  nitric  acid,  ferric  chloride  and  nitric  oxide  are  produced  (p. 
350). 

Ferrous  hydroxide  Fe(OH)2  is  thrown  down  as  a  white  precipitate, 
but  rapidly  becomes  dirty-green  and  finally  brown,  by  oxidation. 
It  dissolves  in  solutions  of  salts  of  ammonium,  being  like  magne- 
sium hydroxide  (p.  525),  sufficiently  soluble  in  water  to  require 
an  appreciable  concentration  of  OH~  for  its  precipitation.  The 
NH4+  from  the  salts  combines  with  the  OH~  formed  by  the  ferrous 
hydroxide  to  give  molecular  ammonium  hydroxide.  Ferrous 
oxide  FeO  is  black,  and  is  formed  by  heating  ferrous  oxalate  in 
absence  of  air.  It  is  made  also  by  cautious  reduction  of  ferric 
oxide  by  hydrogen  (at  about  300°),  but  is  easily  reduced  further 
to  the  metal.  It  catches  fire  spontaneously  when  exposed  to 
the  air. 

Ferrous  carbonate  FeCOs  is  found  in  nature  as  siderite,  and  may 
be  made  in  slightly  hydrolyzed  form  by  precipitation.  The  pre- 
cipitate is  white  but  rapidly  darkens  and  finally  becomes  brown, 
the  ferrous  hydroxide  produced  by  hydrolysis  being  oxidized  to  the 
ferric  condition.  The  salt  interacts  with  water  containing  car- 


634  COLLEGE    CHEMISTRY 

bonic  acid,  after  the  manner  of  calcium  carbonate  (p.  383),  giving 
FeH2(CO3)2,  and  hence  is  found  in  solution  in  natural  (chalybeate) 
waters. 

Ferrous  sulphide  FeS  may  be  formed  as  a  black,  metallic-looking 
mass  by  heating  together  the  free  elements.  It  is  produced  by 
precipitation  with  ammonium  sulphide,  but  not  with  hydrogen  sul- 
phide. It  interacts  readily  with  dilute  acids.  The  precipitated 
form  is  slowly  oxidized  to  ferrous  sulphate  by  the  air. 

Ferrous  sulphate  is  obtained  by  allowing  pyrites  to  oxidize  in  the 
air  and  leaching  the  residue: 

2FeS2  +  702  +  2H20  ->  2FeS04  +  2H2S04. 

The  liquor  is  treated  with  scrap  iron  and  the  neutral  solution  evapo- 
rated until  a  hydrate  FeSO4,7H2O,  green  vitriol,  or  "copperas,"  is 
deposited.  The  crystals  are  efflorescent,  and  become  also  brown 
from  oxidation  to  a  basic  ferric  sulphate: 

4FeS04  +  02  +  2H20-+4Fe(OH)S04. 

With  excess  of  sulphuric  acid  and  air,  or  an  oxidizing  agent  such  as 
nitric  acid,  ferric  sulphate  is  formed.  The  ferrous  sulphate  is  used 
in  dyeing  and  in  making  writing-ink.  The  extract  of  nut-galls  con- 
tains tannic  acid,  HCuHgOg,  which,  with  ferrous  sulphate,  gives 
ferrous  tannate,  a  soluble,  almost  colorless  salt.  A  solution  of  this 
salt  containing  gum-arabic  and  some  blue  or  black  dye  constitutes 
the  ink.  When  the  writing  is  exposed  to  the  air,  the  ferrous 
tannate  is  oxidized  to  the  ferric  condition,  and  the  ferric  compound 
is  a  fine,  black  precipitate  (cf.  p.  516).  The  dye  is  added  to  make 
the  writing  visible  from  the  first.  Ferrous  sulphate  is  also  used  in 
the  purification  of  water  (p.  560). 

Ferric  Compounds.  —  By  leading  chlorine  into  a  solution  of 
ferrous  chloride,  and  evaporating  until  the  proper  proportion 
of  water  alone  remains,  a  yellow,  deliquescent  hexahydrate  of 
ferric  chloride,  FeCl3,6H2O  is  obtained.  When  this  is  heated  still 
further,  hydrolysis  takes  place  and  the  oxide  remains.  When 
chlorine  is  passed  over  heated  iron,  anhydrous  ferric  chloride 
sublimes  in  dark  green  scales,  which  are  red  by  transmitted  light. 
In  solution,  the  salt,  like  other  ferric  salts,  can  be  reduced  to  the 


IRON  635 

ferrous  condition  by  boiling  with  iron.     The  same  reduction  is 
effected  by  hydrogen  sulphide: 


2Fe+++  +  S=->2Fe++  +  S  j. 

The  ferric  ion  is  almost  colorless,  the  yellow-brown  color  of  solu- 
tions of  ferric  chloride  being  due  to  the  presence  of  ferric  hydroxide 
produced  by  hydrolysis.  The  color  deepens  when  the  solution  is 
heated  (increased  hydrolysis),  and  fades  again  very  slowly,  by 
reversal  of  the  action,  when  the  cold  solution  is  allowed  to  stand. 

Ferric  hydroxide  Fe(OH)3  appears  as  a  brown  precipitate  when 
a  base  is  added  to  a  ferric  salt.  It  does  not  interact  with  excess 
of  the  alkali.-  In  this  form  the  substance  dries  to  the  oxide  with- 
out giving  definite  intermediate  hydrated  oxides.  The  hydrates, 
Fe2O3,2Fe(OH)3  (brown  iron  ore)  and  Fe203,4Fe(OH)3  (bog  iron 
ore),  however,  are  found  in  nature  (see  Rust,  p.  632).  The  hy- 
droxide passes  easily  into  colloidal  solution  in  a  solution  of  ferric 
chloride,  and  by  subsequent  dialysis  through  a  piece  of  parchment 
the  salt  can  be  separated,  and  a  pure  colloidal  suspension  of  the 
hydroxide  obtained.  This  suspension,  known  as  dialyzed  iron, 
is  red  in  color,  shows  no  depression  in  the  freezing-point,  and  is 
not  an  electrolyte.  The  hydroxide  is  a  positive  colloid  and  is 
coagulated  (brown  precipitate)  by  the  addition  of  salts,  bivalent 
negative  ions  being  more  effective  than  univalent  ones  (p.  417). 

Ferric  oxide,  Fe2O3,  is  sold  as  "rouge"  and  "  Venetian  red."  It 
is  made  from  the  ferrous  sulphate,  obtained  in  cleaning  iron  ware 
which  is  to  be  tinned  or  galvanized,  and  in  other  ways  in  the  arts. 
The  salt  is  allowed  to  oxidize,  and  the  ferric  hydroxide,  thrown 
down  by  the  addition  of  lime,  is  calcined.  The  product  varies 
in  tint  from  a  bright  yellowish-red  to  a  dark  violet-brown  according 
to  the  fineness  of  the  powder.  The  best  rouge  is  obtained  by 
calcining  ferrous  oxalate  FeC204.  This  oxide  is  not  distinctly 
acidic,  but  by  fusion  with  more  basic  oxides,  compounds  like 
franklinite  Zn(Fe02)2  may  be  formed.  It  is  reduced  by  hydrogen, 
at  about  300°  to  ferrous  oxide,  and  at  700-800°  to  metallic  iron. 

Magnetic  oxide  of  iron  Fe3O4  or  lodestone  is  found  in  nature,  and 
is  formed  by  the  action  of  air  (hammer-scale),  steam,  or  carbon 
dioxide  on  iron.  It  forms  octahedral  crystals,  and  is  a  ferrous- 
ferric  oxide  FeO,Fe2Oa  or  Fe(FeO2)2,  related  to  franklinite. 


636  COLLEGE    CHEMISTRY 

Ferric  sulphide  Fe2S3  may  be  made  by  fusing  together  the  free 
elements,  and  is  obtained  also  as  a  precipitate  by  the  addition 
of  ammonium  sulphide  to  ferric  chloride  solution  (Stokes).  With 
hydrogen  sulphide,  only  sulphur  is  thrown  down  (p.  635). 

Ferric  sulphate  Fe2(SO4)3  is  formed  by  oxidation  of  ferrous  sul- 
phate, and  is  obtained  as  a  white  mass  by  evaporation.  It  gives 
alums  (p.  558),  such  as  ferric-ammonium  alum  (NH4)2S04,Fe2(SO4)3, 
24H2O,  which  are  almost  colorless  when  pure,  but  usually  have  a 
pale  reddish-violet  tinge. 

Pyrite.  —  The  mineral  pyrite  FeS2  (Fools'  gold)  is  the  sulphide 
of  iron  which  is  most  stable  in  the  air.  It  is  found  in  nature  in 
the  form  of  glittering,  golden-yellow  cubes,  octahedrons,  and  pen- 
tagonal dodecahedrons.  It  is  not  attacked  by  dilute  acids,  but 
concentrated  hydrochloric  acid  slowly  converts  it  into  ferrous 
chloride  and  sulphur.  It  is  reduced  by  hydrogen  to  ferrous 
sulphide. 

Cyanides.  —  When  potassium  cyanide  is  added  to  solutions  of 
ferrous  or  ferric  salts,  yellowish  precipitates  are  produced,  but  the 
simple  cyanides  cannot  be  obtained  in  pure  form.  These  precipi- 
tates interact  with  excess  of  the  cyanide  giving  soluble  complex 
cyanides  of  the  forms  4KCN,Fe(CN)2  and  3KCN,Fe(CN)3. 
These  are  called  ferro-  and  ferricyanide  of  potassium,  respectively. 
Ferrocyanide  of  potassium  K4Fe(CN)6,3H20,  "yellow  prussiate 
of  potash/'  is  made  by  heating  nitrogenous  animal  refuse,  such  as 
blood,  with  iron  filings  and  potassium  carbonate.  The  resulting 
mass  contains  potassium  cyanide  and  ferrous  sulphide,  and  when 
it  is  treated  with  warm  water  these  interact  and  produce  the  ferro- 
cyanide : 

2KCN  +  FeS  -»  Fe(CN)2  +  K2S, 
4KCN  +  Fe(CN)2  -»  K4.Fe(CN)6. 

The  salt  is  made  also  from  the  cyanogen  contained  in  crude  illumi- 
nating gas.  The  trihydrate  forms  large,  yellow,  monosymmetric 
tables.  The  solution  contains  almost  exclusively  the  ions  K+  and 
Fe(CN)6— =,  and  gives  none  of  the  reactions  of  the  ferrous  ion  Fe++. 
The  corresponding  acid  H4.Fe(CN)6  may  be  obtained  as  white 
crystalline  scales  by  addition  of  an  acid  and  of  ether  (in  which  the 
substance  is  less  soluble  than  in  water)  to  the  salt.  The  acid  is  a 


IRON  637 

fairly  active  one,  but  is  unstable  and  decomposes  in  a  complex 
manner.  Other  ferrocyanides  may  be  made  by  precipitation. 
That  of  copper  Cu2.Fe(CN)6  is  brown,  and  ferric  ferrocyanide 
Fe4[Fe(CN)6]3  has  a  brilliant  blue  color  (Prussian  blue).  The  fer- 
rous compound  (insoluble)  Fe2Fe(CN)6,  or  perhaps  K2FeFe(CN)6, 
is  white  but  quickly  becomes  blue  by  oxidation.  The  soluble 
ferrocyanides  are  not  poisonous. 

Ferricyanide  of  potassium  K3Fe(CN)6  is  easily  made  from  the 
ferrocyanide  by  oxidation: 

2K4Fe(CN)6  +  C12-+2KC1  +  2K3.Fe(CN)6, 
or  2Fe(CN)6==  +  C12  -*  2Fe(CN)6=-  +  2C1". 

It  forms  red  monosymmetric  prisms.  The  free  acid  H3Fe(CN)6 
is  unstable.  Other  salts  may  be  prepared  by  precipitation. 
Ferrous  ferricyanide  Fe3[Fe(CN6)]2  is  deep-blue  in  color  (Turn- 
bull's  blue).  With  ferric  salts  only  a  brown  solution  is  obtained. 
Prussian  blue  and  Turnbull's  blue  are  used  in  making  laundry 
blueing.  They  are  insoluble,  but  give  colloidal  suspensions  and 
are  adsorbed  by  the  material  of  the  cloth. 

Blue-Prints.  —  Some  ferric  salts,  when  exposed  to  light,  are 
reduced  to  the  ferrous  condition.  Thus,  ferric  oxalate,  in  the 
light,  gives  ferrous  oxalate: 

Fe2(C2O3)3  ->2FeC203  +  2CO2. 

When  paper  is  coated  with  ferrous  oxalate  solution  and  dried, 
and  an  ink  drawing  on  transparent  paper  is  placed  over  the  pre- 
pared surface,  sunlight  will  reduce  the  iron  to  the  ferrous  condi- 
tion, excepting  where  the  ink  protects  it.  When  the  sheet  is  then 
dipped  in  potassium  ferricyanide  solution  (developer),  the  ferric 
oxalate  gives  only  the  brown  substance  which  can  be  washed  out. 
But  the  deep  blue,  insoluble  ferrous  ferricyanide  is  precipitated 
in  the  pores  of  the  paper  where  the  light  has  acted.  The  drawing 
appears  white  on  a  blue  background.  In  ordinary  blue-print 
paper,  ammonium-ferric  citrate  takes  the  place  of  the  oxalate, 
and  the  ferricyanide  has  already  been  applied  to  the  paper  before 
drying,  so  that  only  exposure  and  washing  remain  to  be  done. 
Dilute  sodium  hydroxide  solution  decomposes  the  ferricyanide, 
and  is  used  for  writing  (in  white)  on  blue-prints. 


638  COLLEGE    CHEMISTRY 

Iron  Carbonyls.  —  When  carbon  monoxide  is  led  over  finely 
divided  iron  at  40-80°,  or  under  eight  atmospheres  pressure  at 
the  ordinary  temperature,  volatile  compounds  of  the  composition. 
Fe(CO)4,  iron  tetracarbonyl,  and  Fe(CO)5,  the  pentacarbonyl,  are 
formed.  When  the  gaseous  mixture  is  heated  more  strongly,  the 
compounds  decompose  again,  and  iron  is  deposited.  Illuminating- 
gas  burners  frequently  receive  a  deposit  of  iron  from  this  cause. 

Analytical  Reactions  of  Compounds  of  Iron.  —  There  are 
two  ionic  forms  of  iron,  ferrous-ion  Fe++,  which  is  very  pale-green, 
and  ferric-ion  Fe44^,  which  is  almost  colorless.  Ammonium 
sulphide  gives  with  the  former  black  ferrous  sulphide,  which  is 
soluble  in  dilute  acids.  The  hydroxides  are  white  and  brown, 
respectively,  and  ferrous  carbonate  is  white.  With  ferric  salts, 
which  are  hydrolyzed  (about  5%),  carbonates  yield  the  hydroxide 
because  they  neutralize  the  free  acid  and  displace  the  equilibrium. 
With  ferrocyanide  of  potassium,  ferrous  salts  give  a  white,  and 
ferric  salts  a  blue  precipitate.  With  ferricyanide  of  potassium 
the  former  gives  a  deep-blue  precipitate,  and  the  latter  a  brown 
solution.  Ferric  thiocyanate  Fe(CNS)3  is  deep-red  (p.  182).  With 
borax,  iron  compounds  give  a  bead  which  is  green  (ferrous  borate) 
in  the  reducing  flame,  and  colorless  or,  with  much  iron,  yellow 
(ferric  borate)  or  even  brown  when  oxidized. 

COBALT  Co 

The  Chemical  Relations  of  the  Element.  —  Cobalt  forms 
cobaltous  and  cobaltic  oxides  and  hydroxides  CoO  and  Co(OH)2, 
Co203  and  Co(OH)3,  respectively,  which  are  all  basic,  the  former, 
more  so  than  the  latter.  The  cobaltous  salts  are  little  hydrolyzed, 
but  the  cobaltic  salts  are  largely  decomposed  by  water.  The 
latter  also  liberate  readily  one-third  of  the  negative  radical,  after 
the  manner  of  manganic  salts,  becoming  cobaltous.  Complex 
cations  and  anions  containing  cobalt  are  very  numerous  and  very 
stable. 

Occurrence  and  Properties.  —  Cobalt  is  found  along  with 
nickel  in  smalt ite  CoAs2  and  cobalt ite  CoAsS.  The  pure  metal 
may  be  made  by  Goldschmidt's  process,  or  by  reducing  the  oxalate, 
or  an  oxide,  with  hydrogen. 


COBALT  639 

The  metal  is  silver-white,  with  a  faint  suggestion  of  pink.  It  is 
marked  by  crystalline,  less  tough  than  iron,  and  has  no  commercial 
applications.  It  displaces  hydrogen  slowly  from  dilute  acids,  but 
interacts  readily  with  nitric  acid. 

Cobaltous  Compounds.  —  The  chloride  CoCl2,6H2O  may  be 
made  by  treating  the  oxide  with  hydrochloric  acid.  It  forms  red 
prisms,  and  when  partially  or  completely  dehydrated  becomes 
deep-blue.  Writing  made  with  a  diluted  solution  upon  paper  is 
almost  invisible,  but  becomes  blue  when  warmed  and  afterwards 
takes  up  moisture  from  the  breath,  and  is  once  more  invisible 
(sympathetic  ink).  Most  cobaltous  compounds  are  red  when 
hydrated  or  in  solution  (Co++),  and  blue  when  dehydrated.  By 
addition  of  sodium  hydroxide  to  a  cobaltous  salt,  a  blue  basic 
salt  is  precipitated.  When  the  mixture  is  boiled,  the  pink  cobalt- 
ous hydroxide  Co(OH)2  is  formed.  This  becomes  brown  through 
oxidation  by  the  air.  It  interacts  with  ammonium  hydroxide, 
giving  a  soluble  ammonio-cobaltous  hydroxide,  which  is  quickly 
oxidized  by  the  air  to  an  ammonio-cobaltic  compound  (see  below). 
It  dissolves  also  in  salts  of  ammonium  as  does  magnesium  hy- 
droxide (p.  525).  When  dehydrated  it  leaves  the  black  cobaltous 
oxide  CoO.  Cobaltous  sulphate,  CoSO4,7H2O,  and  cobaltous  ni- 
trate, Co(NO3)2,6H2O,  are  familiar  salts.  The  black  cobaltous 
sulphide  CoS  is  precipitated  by  ammonium  sulphide  from  solu- 
tions of  all  salts,  and  even  by  hydrogen  sulphide  from  the  acetate, 
or  a  solution  containing  much  sodium  acetate  (cf.  p.  484).  Once 
it  has  been  formed,  it  interacts  very  slowly  even  with  dilute  hy- 
drochloric acid,  having  apparently  changed  into  a  less  active 
form.  A  sort  of  cobalt  glass,  made  by  fusing  sand,  cobalt  oxide, 
and  potassium  nitrate,  forms,  when  powdered,  a  blue  pigment, 
smalt,  used  in  china-painting  and  by  artists. 

Cobaltic  Compounds.  —  By  addition  of  a  hypochlorite  to  a 
solution  of  a  cobaltous  salt,  cobaltic  hydroxide  Co(OH)3,  a  black 
powder,  is  precipitated.  Cautious  ignition  of  the  nitrate  gives 
cobaltic  oxide  Co2O3.  Stronger  ignition  gives  the  commercial 
oxide,  which  is  a  cobalto-cobaltic  oxide  CoaC^.  Cobaltic  oxide 
dissolves  in  cold  hydrochloric  acid,  but  the  solution  gives  off 
chlorine  when  warmed.  By  placing  cobaltous  sulphate  solution 


640  COLLEGE    CHEMISTRY 

round  the  anode  of  an  electrolytic  cell,  crystals  of  cobaltic  sul- 
phate, €02(804)3,  have  been  made  and  cobaltic  alums  have  also 
been  prepared  (Hugh  Marshall). 

Complex  Compounds.  —  Potassium  cyanide  precipitates  from 
cobaltous  salts  a  brownish-white  cyanide.  This  interacts  with 
excess  of  the  reagent,  giving  a  solution  of  potassium  cobaltocy- 
anide  K4.Co(CN)6  (cf.  p.  636).  This  compound  is  easily  oxidized 
by  chlorine,  or  even  when  the  solution  is  boiled  in  the  air,  and  the 
colorless  potassium  cobalticyanide  is  formed: 

4K4Co(CN)6  +  2H20  +  O2  -» 4K3.Co(CN)6  +  4KOH. 

The  solution  gives  none  of  the  reactions  of  Co+++,  and  with  acids 
the  very  stable  cobalticyanic  acid,  H3Co(CN)3,  is  liberated. 

When  acetic  acid  and  potassium  nitrite  are  added  to  a  cobaltous 
salt,  the  latter  is  oxidized  by  the  nitrous  acid  (liberated  by  the 
acetic  acid)  and  a  white  complex  salt  K3.Co(NO2)6,nH20  (  = 
Co(N02)3,3KNO2),  potassium  cobaltinitrite,  is  thrown  down. 

Cobaltic  salts  give  with  ammonia  complex  compounds  which  are 
many  and  various.  The  cations  often  contain  negative  groups,  and 
are  such  as  Co(NH3)6+++,  Co(NH3)6Cl+++,  and  Co(NH3)5NO2+++. 
Usually  the  solutions  give  none  of  the  reactions  of  cobaltic  ions, 
and  often  fail  likewise  to  give  those  of  the  anion  of  the  original  salt. 

NICKEL  Ni 

The  Chemical  Relations  of  the  Element.  —  Nickel  forms 
nickelous  and  nickelic  oxides  and  hydroxides  NiO  and  Ni(OH)2, 
Ni203,  and  Ni(OH)3,  but  only  the  former  are  basic.  The  nickel- 
ous salts  resemble  the  cobaltous  and  ferrous  salts,  but  are  not 
oxidizable  into  corresponding  nickelic  compounds.  Since  there 
are  no  nickelic  salts,  there  are  here  no  analogues  of  the  cobalti- 
cyanides  or  the  cobaltinitrites.  The  complex  nickelous  salts, 
like  the  complex  cobaltous  salts,  and  unlike  the  complex  cobaltic 
salts,  are  unstable,  and  so  give  some  of  the  reactions  of  Ni++. 

Occurrence  and  Properties.  —  Nickel  occurs  free  in  meteor- 
ites and  in  niccolite  NiAs  and  nickel  glance  NiAsS.  It  is  now 
manufactured  chiefly  from  pentlandite  [Ni,Cu,Fe]S  and  other 


NICKEL  641 

minerals  found  at  Sudbury  (Ontario),  and  from  garnierite,  a 
silicate  of  nickel  and  magnesium,  found  in  New  Caledonia.  In 
the  former  case,  the  ore  is  roasted,  smelted,  and  finally  bessem- 
erized.  The  resulting  alloy  of  copper  and  nickel  is  much  used 
for  sheet-metal  work  (Monel  metal,  approx.  1  :  1).  Pure  nickel 
is  separated  from  the  copper  by  an  electrolytic  process  (p.  511), 
or  by  the  Monde  process  (see  below). 

The  metal  is  white,  with  a  faint  tinge  of  yellow,  is  very  hard,  and 
takes  a  high  polish  (m.-p.  1452°).  It  is  used  in  making  alloys, 
such  as  German  silver  (copper,  zinc,  nickel,  2:1:1)  and  the 
" nickel"  used  in  coinage  (copper,  nickel,  3:1).  Although  in 
these  alloys  the  red  color  of  the  copper  is  completely  lost,  the 
copper  is  simply  dissolved,  and  not  combined.  Zinc  and  copper, 
however,  give  a  compound  Cu2Zn3.  Nickel  plating  on  iron  is 
accomplished  exactly  like  silver  plating  (p.  516).  The  bath  con- 
tains an  ammoniacal  solution  of  ammonium-nickel  sulphate 
(NH4)2S04,NiSO4,6H20,  and  a  plate  of  nickel  forms  the  anode. 

The  metal  rusts  very  slowly  in  moist  air.  It  displaces  hydro- 
gen with  difficulty  from  dilute  acids;  but  interacts  with  nitric 
acid. 

Compounds  of  Nickel.  —  The  chloride  NiCl2,6H20  is  made  by 
treating  any  of  the  oxides  with  hydrochloric  acid,  and  is  green 
in  color  (when  anhydrous,  brown).  The  sulphate  NiS04,6H2O, 
which  crystallizes  in  green,  square  prismatic  forms  at  30-40°,  is 
the  most  familiar  salt.  Nickelous  hydroxide,  Ni(OH)2,  is  formed 
as  an  apple-green -precipitate,  and  when  heated  leaves  the  green 
nickelous  oxide  NiO.  It  dissolves  in  ammonium  hydroxide, 
giving  a  complex  nickel-ammonia  cation.  It  is  soluble  also  in 
salts  of  ammonium  (c/,  p.  525).  By  cautious  ignition  of  the 
nitrate,  nickelic  oxide  Ni2O3  is  formed  as  a  black  powder.  The 
oxides  and  salts,  when  heated  strongly  in  oxygen,  give  the  oxide 
Ni304.  The  last  two  oxides  liberate  chlorine  when  treated  with 
hydrochloric  acid,  and  give  nickelous  chloride.  Nickelic  hydrox- 
ide Ni(OH)3  is  a  black  precipitate  formed  when  a  hypochlorite 
is  added  to  any  salt  of  nickel.  Nickelous  sulphide  is  thrown  down 
by  ammonium  sulphide,  and  behaves  like  cobaltous  sulphide 
(p.  639).  It  forms  a  brown  colloidal  solution  when  excess  of  the 
precipitant  is  used,  and  is  then  deposited  very  slowly. 


642  COLLEGE   CHEMISTRY 

Addition  of  dimethylglyoxime  to  an  ammoniacal  solution  of  a 
salt  of  nickel  gives  a  brilliant  scarlet  precipitate  of  an  acid  salt: 

Ni(OH)2  +  2(HON)2C2(CH3)2-+2H20  +  NiH2[C2N202(CH3)2]2. 

This  reaction  is  not  shown  by  salts  of  cobalt,  especially  if  oxidation 
to  the  cobaltic  condition  has  been  permitted  by  contact  with  air. 

With  potassium  cyanide  anql  a  salt  of  nickel  the  greenish  nickel- 
ous  cyanide,  Ni(CN)2,  is  first  precipitated.  This  dissolves  in 
excess  of  the  reagent,  and  a  complex  salt  K2Ni(CN)4,H2O  (  = 
2KCN.Ni(CN)2)  may  be  obtained  from  the  solution.  This  salt 
is  of  different  composition  fn  "n  the  corresponding  compounds  of 
cobalt  and  of  iron,  and  is  stable.  Thus,  with  bleaching 

powder,  it  gives  Ni(OH)3  as  i  ..ack  precipitate.  When  the  solu- 
tion is  boiled  in  the  air  no  oxidation  to  a  complex  nickelicyanide 
occurs,  and  indeed  no  such  salts  are  known.  This  fact  enables 
the  chemist  to  separate  col)..  and  nickel,  for  when  the  mixed 
cyanides  are  boiled  and  then  reated  with  bleaching  powder,  the 
cobalticyanide  is  unaffected.  With  potassium  nitrite  and  acetic 
acid  no  insoluble  compound  corresponding  to  that  given  by  cobalt 
salts  is  formed  by  salts  of  >  ickel.  The  only  known  compound 
which  could  be  formed,  4KN<O2,Ni(N02)2,  is  soluble.  This  action 
also  is  used  for  the  purpose  of  separation.  The  pink  color  of 
cobalt  salts  and  the  green  of  nickel  salts  are  complementary  colors, 
so  that,  by  using  suitable  proportions  of  the  two,  a  colorless  mix- 
ture can  be  produced. 

When  finely  divided  nickel,  made  by  reducing  the  oxide  or 
oxalate  with  hydrogen  at  a  moderate  temperature,  is  exposed  to  a 
stream  of  cold  carbon  monoxide,  nickel  carbonyl  Ni(CO)4  is  formed. 
This  is  a  vapor  and  is  condensable  to  a  colorless  liquid  (b.-p.  43° 
and  m.-p.  —25°).  The  vapor  is  poisonous.  When  heated  to 
150-180°  it  is  dissociated  and  nickel  is  deposited.  Cobalt  forms 
no  corresponding  compound.  Commercially,  pure  nickel  is  sepa- 
rated from  copper  (and  cobalt)  in  the  Monde  process  by  passing 
carbon  monoxide  over  the  pulverized  alloy,  and  subsequently 
heating  the  gas. 

Analytical  Reactions  of  Compounds  of  Cobalt  and  Nickel. 

—  The  cobalt  ion  Co++  is  pink,  and  the  nickelous  ion  Ni++  green. 
The  reactions  used  in  analysis  have  been  described  in  the  preceding 


NICKEL  643 

paragraphs.  With  borax,  cobalt  compounds  give  a  blue  bead 
(cobaltous  borate),  and  nickel  compounds  a  bead  which  is  brown 
in  the  oxidizing  flame  and  cloudy,  from  the  presence  of  gray, 
metallic  nickel,  when  reduced. 

Exercises.  —  1.  What  would  be  the  interactions  of  calcium  car- 
bonate when  fused  with  sand  and  with  clay,  respectively? 

2.  Make  equations  representing,   (a)  the  oxidation  of  ferrous 
chloride  by  air,  (6)  the  hydrolysis  of  ferrous  carbonate  and  the 
oxidation  of  ferrous  hydroxide,  (c)  'the  oxidation  of  ferrous  sul- 
phate with  excess  of  sulphuric  acid  by  hypochlorous  acid,   (d) 
the  formation  of  ferrous  and  ferric  tannates  (p.  634),  (e)  the  re- 
duction of  ferric  chloride  by  iron*  and  by  hydrogen  sulphide, 
respectively,  (/)  the  dry  distillation  of  basic  ferric  sulphate,  (g) 
the  formation  of  ferric  ferrocyanide  and  of  ferrous  ferricyanide. 

3.  Explain  the  solubility  of  cob    '.ous  and  nickelous  hydroxide 
in  salts  of  ammonium. 

4.  Construct  equations  to  show  the  formation,  (a)  of  the  in- 
soluble potassium  cobaltinitrite  (nitric  oxide  is  given  off),  (b)  of 
nickelic   hydroxide   from  nickelous   chloride  and  sodium  hypo- 
chlorite.     Remembering  that  the  hyj.  ochlorite  is  somewhat  hydro- 
lyzed,  explain  why  the  precipitation  in  (6)  is  complete. 

5.  Tabulate  in  detail  the  chemical  relations  of  the  elements 
cobalt  and  nickel,  with  especial  reference  to  showing  the  resem- 
blances and  differences. 


CHAPTER  XLVI 
THE   PLATINUM   METALS 

THE  rarer  elements  of  Mendelejeff's  eighth  group  divide  them- 
selves into  sets  of  three  each.  Just  as  iron,  cobalt,  and  nickel 
have  similar  atomic  weights  and  much  the  same,  specific  gravity 
(7.8-8.8),  so  ruthenium  (Ru,  at.  wt.  101.7),  rhodium  (Rh,  at.  wt 
103),  and  palladium  (Pd,  at.  wt.  106.7)  have  specific  gravities 
from  12.26  to  11.5.  Similarly  osmium  (Os,  at.  wt.  191),  indium 
(Ir,  at.  wt.  193),  and  platinum  (Pt,  at.  wt.  195.2)  form  a  triad 
with  specific  gravities  from  22.5  to  21.5.  Chemically,  ruthenium 
shows  the  closest  resemblance  to  osmium,  and  both  are  allied  tc 
iron.  Similarly,  rhodium  and  iridium,  and  palladium  and  plati- 
num are  natural  pairs. 

The  six  elements  are  found  alloyed  in  nuggets  and  particles 
which  are  separated  from  alluvial  sand  by  washing.  Platinun: 
forms  60-^4  per  cent  of  the  whole.  The  chief  deposits  are  in  the 
Ural  Mountains,  smaller  amounts  being  found  in  California 
Australia,  Borneo,  and  elsewhere.  The  components  are  separates 
by  a  complex  series  of  chemical  operations. 

Ruthenium  and  Osmium.  —  These  metals  are  gray  like  iron 
while  the  other  four  are  whiter  and  more  like  cobalt  and  nickel 
They  also  resemble  iron  in  being  the  most  infusible  members  oi 
their  respective  sets.  Both  melt  considerably  above  2000° 
They  likewise  resemble  iron  in  uniting  easily  with  free  oxygen 
while  the  other  four  elements  do  not.  Ruthenium  gives  RuO: 
and  even  RuO4,  although  the  latter  oxide  is  more  easily  obtainec 
indirectly.  Osmium  gives  Os04,  "osmic  acid,"  a  white  crystal- 
line body  melting  at  40°  and  boiling  at  about  100°.  The  odor 
and  irritating  effects  of  the  vapor  recall  chlorine  (Gk.  007*77,  odor), 
The  substance  is  not  an  acid,  nor  even  an  acid  anhydride.  The 
aqueous  solution  is  used  in  histology,  and  stains  tissues  in  conse- 
quence of  its  reduction  by  organic  bodies  to  metallic  osmium, 
It  is  affected  particularly  by  fat.  Osmic  acid  also  hardens  the 

644 


THE    PLATINUM    METALS  645 

material  without  distorting  it.  Osmium  forms  also  a  yellow, 
crystalline  fluoride,  OsF8  (m.-p.  34.5°).  It  will  be  observed  that 
ruthenium  and  osmium  have  a  maximum  valence  of  eight. 

Rhodium  and  Iridium.  —  These  metals  are  not  attacked  by 
aqua  regia,  while  the  other  four  are  dissolved,  more  or  less  slowly. 
They  are  harder  than  platinum,  and  iridium  is  alloyed  with  this 
metal  for  the  purpose  of  increasing  its  resistance  to  the  action  of 
acids.  They  resemble  cobalt  in  having  no  acid-forming  properties. 
The  most  familiar  compounds  of  iridium  are  the  complex  chlorides 
X3IrCl6  (=  3XCl,IrCl3)  and  X2IrCl6  (=  2XCl,IrCl4).  The  solu- 
tions of  the  latter  are  red,  and  the  acid,  chloro-iridic  acid  H2IrCls, 
is  often  found  in  commercial  chloroplatinic  acid  H^PtCle,  and 
confers  upon  it  a  deeper  color. 

Palladium  and  Platinum.  —  Palladium  is  the  only  metal  of 
this  family  which  is  attacked  by  nitric  acid.  Palladium  and  plati- 
num form  -ous  and  -ic  compounds  of  the  forms  PdX2  and  PdXi, 
respectively.  The  oxides  PdO  and  PtO  and  corresponding  hy- 
droxides are  basic.  When  quadrivalent,  the  metals  appear  chiefly 
in  complex  compounds,  like  H2.PtCle  and  H2.PdCle,  in  which  the 
metal  is  in  the  anion.  Platinum  gives  also  platinates  derived  from 
the  oxide  Pt02. 

Palladium.  —  This  metal  (m.-p.  1549°),  named  from  the  planet- 
oid Pallas,  is  noted  chiefly  for  its  great  tendency  to  absorb  hy- 
drogen. When  finely  divided,  it  takes  up  about  800  times  its 
own  volume.  The  amount  absorbed  varies  continuously  with 
the  concentration  (pressure)  of  the  hydrogen,  although  not  ac- 
cording to  a  uniform  rule,  and  the  product  is  in  part  at  least  a 
solid  solution.  When  a  strip  of  palladium  is  made  the  cathode 
of  an  electrolytic  cell,  over  900  volumes  of  hydrogen  may  be 
occluded.  This  absorbed  hydrogen,  in  consequence  of  the  cata- 
lytic influence  of  the  metal,  reacts  more  rapidly  than  does  the  gas, 
and  consequently  a  strip  of  hydrogenized  palladium  will  quickly 
precipitate,  from  solutions  of  their  salts,  copper  and  other  metals 
less  electropositive  than  hydrogen  and  will  reduce  ferric  and  other 
reducible  salts: 

CuS04  +  H2  -»  H2S04  +  Cu,      or      Cu++  +  H2  -»  2H+  +  Cu. 
2FeCl3  +  H2  ->  2FeCl2  +  2HC1,  or  2Fe+++  +  H2  ->  2Fe++  +  2H+. 


646  COLLEGE    CHEMISTRY 

Platinum.  —  This  metal  (dim.  of  Sp.  plata,  silver)  is  grayish- 
white  in  color,  and  is  very  ductile.  At  a  red  heat  it  can  be  welded. 
It  does  not  melt  in  the  Bunsen  flame,  but  fuses  easily  in  the  oxyhy- 
drogen  jet  (m.-p.  1755°).  On  account  of  its  very  small  chemical 
activity  it  is  used  in  electrical  apparatus  and  for  making  wire, 
foil,  and  crucibles  and  other  vessels  for  use  in  laboratories.  It 
interacts  with  fused  alkalies,  giving  platinates.  The  oxygen  acids 
are  without  action  upon  it,  but  on  account  of  the  tendency  to 
form  the  extremely  stable  complex  ion  *PtCle=  (p.  520),  the  free 
chlorine  and  chloride-ion  in  aqua  regia  convert  it  into  chloro- 
platinic  acid  H^PtCle. 

The  metal  condenses  oxygen  upon  its  surface  and  it  dissolves 
hydrogen.  The  finely  divided  forms  of  the  metal,  such  as  platinum 
sponge  made  by  igniting  ammonium  chloroplatinate  (NH4)2PtCle, 
platinum  black  made  by  adding  zinc  to  chloroplatinic  acid,  and 
platinized  asbestos  made  by  dipping  asbestos  in  a  solution  of  chlo- 
roplatinic acid  and  heating  it,  show  this  behavior  very  conspicu- 
ously. They  cause  instant  explosion  of  a  mixture  of  oxygen  and 
hydrogen,  in  consequence  of  the  heat  developed  by  the  rapid 
union  of  that  part  of  the  gases  which  is  condensed  in  the  metal. 
A  heated  spiral  of  fine  platinum  wire  will  continue  to  glow  if  im- 
mersed in  the  mixture  of  methyl  alcohol  vapor  and  air  (oxygen), 
formed  by  placing  a  little  of  the  alcohol  in  the  bottom  of  a  beaker. 
Some  cigar-lighters  work  on  this  principle.  The  heat  is  developed 
by  the  interaction  between  the  substances,  which  takes  place  with 
great  speed  at  the  surface  of  the  platinum.  Platinum  sponge  is 
used  as  a  contact  agent  in  making  sulphur  trioxide  (p.  279). 

Platinum  was  the  only  otherwise  suitable  substance  which  had  the 
same  coefficient  of  expansion  as  glass,  and  it  was  consequently  fused 
into  incandescent  bulbs  and  furnished  the  electrical  connection 
with  the  filament  in  the  interior.  Recently,  however,  a  less  ex- 
pensive substitute  has  been  found.  Large  amounts  are  also  con- 
sumed in  photography  and  by  dentists.  It  is  used  also  in  making 
jewelry,  and  in  Russia  for  coinage.  The  price  of  the  metal  is 
subject  to  great  variations,  since  a  rainy  season  in  the  Caucasus 
will  render  larger  amounts  accessible  to  the  miners;  but,  on  the 
whole,  the  many  applications  which  have  been  found  for  it  have 
quintupled  its  price  in  the  last  thirty  years.  The  price  is  now 
about  twice  that  of  gold. 


THE    PLATINUM   METALS  647 

When  special  resistance  to  chemical  or  mechanical  influences  is 
required,  as  in  standard  meters  for  international  reference,  or 
points  of  fountain  pens,  the  alloy  with  iridium  is  employed. 

Compounds  of  Platinum.  —  Platinous  chloride  is  made  by 
passing  chlorine  over  finely  divided  platinum  at  240-250°,  or  by 
heating  chloroplatinic  acid  to  the  same  temperature.  It  is  greenish 
and  insoluble  in  water,  but  forms  with  hydrochloric  acid  the 
soluble  chloroplatinous  acid  H2PtCl4.  Potassium  chloroplatinite 
K2PtCl4  is  used  in  making  platinum  prints.  Bases  precipitate 
black  platinous  hydroxide  Pt(OH)2,  which  interacts  with  acids  but 
not  with  bases.  Gentle  heating  gives  the  oxide  PtO  and  stronger 
heating  the  metal.  With  potassium  cyanide  and  barium  cyanide 
soluble  platino-cyanides,  K2Pt(CN)4,3H20  and  BaPt(CN)4,4H20, 
are  formed.  These  substances,  when  solid,  show  strong  fluores- 
cence (p.  606),  converting  X-rays  as  well  as  ultra-violet  rays  into 
visible  radiations.  The  barium  salt  is  used  to  coat  screens  on 
which  the  shadows  cast  by  X-rays  are  received. 

Chloroplatinic  acid  H2PtCl6,6H20  is  made  by  treating  the  metal 
with  aqua  regia,  and  forms  reddish-brown  deliquescent  crystals. 
With  potassium  and  ammonium  salts,  it  yields  the  sparingly  solu- 
ble, yellow  chloroplatinates  K2PtCl6  and  (NH4)2PtCl6  (cf.  p.  452), 
in  solutions  of  which  the  platinum  migrates  towards  the  anode  and 
silver  salts  precipitate  Ag2PtCl6  and  not  silver  chloride.  Platinic 
chloride  PtCl4  is  made  by  heating  chloroplatinic  acid  in  a  stream 
of  chlorine  at  360°.  When  dissolved  in  water,  it  combines  to 
form  H2.PtCl40,  with  the  platinum  in  the  negative  ion.  Bases 
interact  with  chloroplatinic  acid,  giving  a  yellow  or  brown  pre- 
cipitate of  platinic  hydroxide  Pt(OH)4.  This  substance  interacts 
with  bases  to  give  platinates,  like  Na2Hi0Pt3Oi2,H20.  Both  sets 
of  platinum  compounds  interact  with  hydrogen  sulphide,  giving 
the  sulphides  PtS  and  PtS2,  respectively.  These  are  black  powders 
which  dissolve  in  yellow  ammonium  sulphide  solution,  much 
as  do  the  sulphides  of  gold,  arsenic,  and  other  metals,  giving  am- 
monium sulphoplatinates. 


APPENDIX 
I.  The  Metric  System 

Length.  1  meter  (1  m.)  =  10  decimeters  =  100  centimeters  (100 
cm.)  =  1000  millimeters  (1000  mm.). 

1  kilometer  =  1000  meters  (1000  m.)  =  0.6214  miles. 
1  decimeter  =  0.1  m.  =  10  centimeters  =  3.937  inches. 
1  meter  =  1.094  yards  =  3.286  ft.  =  39.37  in. 

Volume.  1  liter  =  1000  cubic  centimeters  (1000  c.c.)  =  a 
cube  10  cm.  X  10  cm.  X  10  cm. 

1  liter  =  0.3532  cu.  ft.  =  61.03  cu.  in.  =  1.057  quarts  (U.  S.)  or 
1.136  quarts  (Brit.)  =  34.1  fl.  oz.  (U/S.)  =  35.3  oz.  (Brit.). 

1  fl.  ounce  (U.  S.)  =  29.57  c.c.     1  ounce  (Brit.)  =  28.4  c.c. 
1  cu.  ft.  =  28.32  liters. 

Weight.  1  gram  (g.)  =  wt.  of  1  c.c.  of  water  at  4°  C.  1  kilo- 
gram =  1000  g. 

1  gram  =  10  decigrams  =  100  centigrams  (100  cgm.)  =  1000 
milligrams  (1000  mgm.). 

1  kilogram  =  2.205  Ibs.  avoird.  (U.  S.  and  Brit.). 
1  Ib.  avoird.  =  453.6  g. 

1  oz.  avoird.  (U.  S.  and  Brit.)  =  28.35  g.     100  g.  =  3.5  oz. 
1  nickel  (U.  S.)  weighs  5  g.     1  halfpenny  (Brit.)  weighs 
5  to  5.7  g. 

II.  Scale  of  Hardness 

Each  of  the  following  minerals  will  scratch  the  surface  of  a 
specimen  of  any  one  preceding  it  in  the  list. 

1.  Talc  6.   Felspar 

2.  Gypsum  (or  NaCl)  7.   Quartz 

3.  Calcite  (or  Cu)  8.   Topaz 

.    4.    Fluor  it  e  9.   Corundum 

5.   Apatite  10.   Diamond 

648 


APPENDIX 


649 


Glass  is  slightly  scratched  by  5,  and  easily  by  those  following. 
Glass  will  not  scratch  5  distinctly,  but  will  scratch  those  preceding  5. 
A  good  knife  scratches  6  slightly,  but  not  those  following. 
A  file  will  scratch  7,  but  not  those  following. 


III.  Temperatures  Centigrade  and  Fahrenheit 

Upon  the  centigrade  scale,  the  freezing-point  of  water  is  0°  C. 
and  the  boiling-point  100°  C.  Upon  the  Fahrenheit  scale,  the 
same  points  are  32°  F.  and  212°  F.,  respectively.  The  same  inter- 
val is  thus  100°  on  the  one  scale  and  180°  on  the  other.  The  degree 
Fahrenheit  is  therefore  j|§  or  |  of  1°  Centigrade.  Any  tempera- 
tures can  be  converted  by  using  the  formulae: 

C.°  =  |  (F.°  -  32),        F.°  =  |  (C.°)  +  32. 

The  following  table  (IV)  contains  the  temperatures  from  0°  C. 
to  35°  C.,  with  the  corresponding  values  on  the  Fahrenheit  scale 
(32°  F.  to95°F.). 


IV.  Vapor  Pressures  of  Water 

Both  the  Fahrenheit  (F.)  or  ordinary  and  the  Centigrade  (C.)  temperatures  are  given. 


Temperature. 

Pressure,  mm. 

Temperature. 

Pressure,  mm. 

F. 

C. 

F. 

C. 

32° 

0° 

4.6 

71.6° 

22° 

19.7 

41 

5 

6.5 

73.4 

23 

20.9 

46.4 

8 

8.0 

75.2 

24 

22.2 

48.2 

9 

8.6 

77.0 

25 

23.6 

50.0 

10 

9.2 

78.8 

26 

25.1 

51.8 

11 

9.8 

80.6 

27 

26.5 

53.6 

12 

10.5 

82.4 

28 

28.1 

55  4 

13 

11.2 

84.2 

29 

29.8 

57.2 

14 

11.9 

86.0 

30 

31.5 

59.0 

15 

12.7 

87.8 

31 

33.4 

60.8 

16 

13.5 

89.6 

32 

35.4 

62.6 

17 

14.4 

91.4 

33 

37.4 

64.4 

18 

15.4 

93.2 

34 

39.6 

66.2 

19 

16.3 

95.0 

35 

41.8 

68  0 

20 

17.4 



...  • 

69.8 

21 

18.5 

212.0 

100 

760^6 

650  COLLEGE    CHEMISTRY 

V.  Order  of  Activity  of  the  Metala 
(Electromotive  Series) 

Each  metal,  when  placed  in  a  solution  of  a  salt  of  one  of  the 
metals  following  it  in  the  list,  displaces  the  second  metal  and 
deposits  it  in  the  free  condition  (see  pp.  60,  260,  438,  531). 

For  explanation  of  potential  differences  (electromotive  series), 
see  pp.  539-547. 

Potassium  Manganese  Tin  Mercury 

Sodium  Zinc  Lead  Silver 

Barium  Chromium  Hydrogen  Palladium 

Strontium  Cadmium  Copper  Platinum 

Calcium  Iron  Arsenic  Gold 

Magnesium  Cobalt      •  Bismuth 

Aluminium  Nickel  Antimony 


INDEX 


***  Acids  are  all  listed  under  "  acid  "  and  salts  under  the  positive  radical. 


ACETONE,  394,  408 
Acetylene,  378,  392,  394,  400 
formula  of,  109 
torch,  394 

Acid,  acetic,  407,  467 
antimonic,  589 
arsenic,  585 
arsenious,  586 
boracic,  431 
boric,  431 
bromic,  318 
carbolic,  349 
carbonic,  383 
chlorauric,  356 
chloric,  314 

chloroplatinic,  356,  647 
chlorous,  314 
chromic,  597 
disulphuric,  285 
formic,  385,  412 
hydrazoic,  340,  345 
hydriodic,  202 
hydrobromic,  198 
hydrochloric,  141 
hydrochloric,  properties,  146 
hydrocyanic,  420 
hydrofluoboric,  431 
hydrofluoric,  206 
hydrofluosilicic,  427 
hydrosulphuric,  269 
hypochlorous,  161,  307,  309 
hyponitrous,  357 
iodic,  318 

metaphosphoric,  368,  371,  417 
metastannic,  569,  571 
nitric,  347 

fuming,  348 

graphic  formula,  358 

oxidizing  actions,  354 

synthetic,  352 

test,  351 

nitrosylsulphuric,  281 
nitrous,  356 

orthophosphoric,  368,  370 
osmic,  644 


Acid,  oxalic,  385,  413 

palmitic,  412 

perchloric,  314,  315 

perchromic,  224 

permanganic,  622 

persulphuric,  291 

phosphoric,  368 

phosphorous,  372 

picric,  349 

pyrophosphoric,  368,  371 

prussic,  420 

selenic,  294 

silicic,  428 

a-stannic,  570 

sulphuric,  279,  280 
graphic  formula,  291 
properties,  285 

sulphurous,  288 

tannic,  634 

thiosulphuric,  290 
Acidic  oxides,  94 
Acidimetry,  255 
Acids,  52,  94 

and  anhydrides,  316,  369 

fractions  ionized,  241 

non-ionic  formation,  261 

of  constant  boiling-point,  145 

organic,  412 

properties  in  solution,  210 
Actinium,  614 
Actions,  non-ionic,  260 

reversible,  177 
Activity,  acids,  242 

apparent,  180 

bases,  242 

chemical,  38,  172 

of  ionogens,  242 

order  of,  metals,  59 

non-metals,  548 
Adsorption,  408,  419 
Affinity,  chemical,  180 
Agate,  427 
Air,  a  mixture,  333 

components  of,  328,  333 

liquid,  334 


651 


652 


INDEX 


Air,  water  vapor  in,  88 

weight  of  22.4  1.,  101 
Alabaster,  485 
Alcohol,  denatured,  407 

ethyl,  406,  407 

methyl,  408 
Alcohols,  413 
Alkalimetry,  255 
Allotropic  modifications,  222 
Alloys,  435,  503,  591 

acid-resisting,  596 

anti-friction,  588 
Alum,  82,  558 

chrome,  603 
Aluminates,  557 
Aluminium,  554 

carbide,  391 

compounds,  556 
Aluminothermy,  556 
Alundum,  558 
Amalgam,  sodium,  345 
Amalgams,  435 
Amethyst,  427 
Ammonia,  340 

household,  344 

properties,  342 

-soda  process,  461 

Ammonio-copper  salts,  504,  506,  507, 
509 

-silver  salts,  515 
Ammonium  amalgam,  455 

carbonate.  211 

compounds,  test  for,  345 

cyanate,  421 

hydroxide,  344 

molybdate,  605 

nitrate,  357 

nitrite,  338 

salts  of,  345,  453 

sulpharsenate,  587 

sulphides,  454 

sulphostannate,  572 

thiocyanate,  421 
Ammono-compounds,  535 
Amorphous  bodies,  97 
Ampere,  137 
Amylase,  406 
Analysis,  qualitative,  537 

volumetric,  257 
Analytical  reactions,  aluminium,  565 

ammonium,  455 

arsenic  family,  593 

cadmium,  531 

calcium,  495 

calcium  family,  498 

***  Acids  are  all  listed  under  "acid 


Analytical  reactions,  chromium,  604 

cobalt  and  nickel,  642 

copper,  510 

iron,  638 

lead,  580 

magnesium,  526 

manganese,  623 

mercury,  536 

potassium,  453 

silver,  518 

sodium,  465 

tin,  572 

zinc.  530 
Anhydride,  and  acid,  369 

and  acid  or  salt,  316 

chromic,  598,  601 

permanganic,  622 
Anhydrides,  94 
Anions,  237 
Anode,  237 
Anthracene,  411 
Antimony,  587 

compounds,  588 
Apatite,  362,  486 
Aq,  52 

Aqua  regia,  356 
Aqueous  tension,  87,  649 

correction  for,  73 

hydrates,  96 
Argentic,  see  silver 
Argon,  335 
Arsenic,  582 

white,  585 
Arsine,  583 
Asphalt,  391 
Assaying,  521 
Atmosphere,  328 
Atom,  constitution,  304 
Atomic  numbers,  303 

weight  of  a  new  element,  118 

weights,  41,  103,  inside  rear  cover 

advantages  of,  107 
Atoms,  43 
Attributes,  19 
Avogadro,  77 

B.T.U.,  409 
Babbitt's  metal,  588 
Baking  powders,  463 

soda,  463 
Barium,  496 

peroxide,  222 
Barometer,  71 
Bases,  94 

fractions  ionized,  242 
and  salts  under  the  positive  radical. 


INDEX 


653 


Bases,  properties  in  solution,  211 

Basic  oxides,  94 

Batteries,  see  cells 

Bead  tests,  372,  433 

Beer,  406 

Benzene,  392,  411 

Benzine,  391 

Beryllium,  523 

Bessemer  process,  629 

Bicarbonates,  383 

Birkeland-Eyde  process,  353 

Bismuth,  591 

Black-lead,  378 

Blast  lamp,  398 

furnace,  626 
Blau  gas,  396 
Bleaching,  311 

hydrogen  peroxide,  224 

powder,  309,  312,  484 
Blue-prints,  637 
Blue-stone,  95,  509 
Body,  definition  of,  4 
Boiling-point,  acid  of  constant,  145 

solutions,  216 
Bone  black,  408    . 
Borax,  431,  432,  465 
Bordeaux  mixture,  509 
Boron,  430 
Brass,  503 
Bromine,  193 

oxygen  acids,  318 

properties,  195 
Bronze,  503, 

Brownian  movement,  416 
Bunsen  flame,  398 
Burette,  256 
Butter,  414 
By-product  coke,  340,  411 

CADMIUM,  530 
Caesium,  453 
Calcining,  275 
Calcite,  83 
Calcium,  474 

bicarbonate,  384 

bisulphite,  290,  402 

carbide,  379,  394 

carbonate,  476 

chloride,  475 

cyanamide,  487 

fluoride,  204,  475 

hydroxide,  477 

hydride,  475 

light,  58 

nitrate,  353 


Calcium,  oxalate,  478-484 

oxide,  477 

phosphate,  362,  486 

phosphide,  365 

silicate,  493 

sulphate,  84,  485 

sulphide,  486 
Calculations,  formulae  from  data,  45 

involving  weights,  66 

volumes,  115,  149 
Calomel,  533 
Calorie,  85 
Calorimeter,  174 
Camphor,  291 
Caramel,  405 
Carbides,  378,  379 
Carbohydrates,  402 
Carbon,  375 

dioxide,  381 

as  plant  food,  387 
uses,  385 

disulphide,  379 

gas,  395 

monoxide,  385 

prints,  600 

tetrachloride,  379,  392 
Carbona,  379 
Carbonates,  383 
Carbonyl  chloride,  163,  387 
Carborundum,  380 
Carnallite,  194,  445 
Catalysts,  definition,  29 

negative,  288 
Cathode,  237 
Cations,  237 

recognition  of,  537 
Cause,  173 
Cell,  Clark,  548 

combination,  540,  541 

concentration,  541,  550 

displacement,  540,  544,  547 

Edison,  579 

gravity,  547 

oxidation,  541,  545 

potential  differences,  546 

storage,  577 

Weston,  548 
CeUuloid,  359 
Cellulose,  402 
Cement,  562 
Cerium,  580 

oxide,  397 
Chalk,  476 
Charcoal,  408 

as  adsorbent,  419 


***  Acids  are  all  listed  under  "acid"  and  salts  under  the  positive  radical. 


654 


INDEX 


Chamber  process,  281 
Chemical  change,  complete,  178 

energy  and,  167 

reversible,  177 

speed  of,  173 

varieties,  7,  14,  55,  147,  148,  166, 
197 

ionic,  259,  270,  504 
Chemical  changes,   concurrent,   315, 
317 

consecutive,  289 
Chemical  equilibrium,  177 

applications,  184 

displacement  of,  185,  203 

history,  187 

in  ammonia,  343 

temperature  and,  188 
Chemical  relations,  163,  192 

halogens,  319 

sulphur  family,  295 
Chlorates,  313 
Chlorides,  preparation,  146 

solubilities,  164 
Chlorine,  154 

dioxide,  314 

monoxide,  307 

not  a  bleacher,  312 

oxides  and  acids,  306 

properties,  159 

-water,  161,  310 
Chloroform,  392 
Chromic  anhydride,  598,  601 

compounds,  602 
Chromite,  597 
Chromium,  595 
Chromous  compounds,  604 
Chromyl  chloride,  601 
Clarke,  F.  W.,  17 
Clay,  430,  561 
Coal,  409 

analysis,  409 

calorific  power,  409 
Coagulation  treatment,  91 
Cobalt,  638 

compounds,  639 
Coke,  411 

by-product  ovens,  340 ' 
Colemanite,  431 
Collie,  15 
Collodion,  359 
Colloids,  403,  415-420,  562 

arsenious  sulphide,  586 

ferric  hydroxide,  635 
Columbium,  594 
Combination,  7 

***  Acids  are  all  listed  under  "acid1 


Combihing     proportions,     measure- 
ment, 33 
Combustion,  35 

spontaneous,  37 
Complex  salts,  505 
Components,  4 
Composition,  definition  of,  34 
Concentration,  chemical  equilibrium 
and,  181 

gases,  76 

Concurrent  reactions,  315,  317 
Conditions,  20 

Conductivity  of  ionogens,  239 
Congo  red,  258,  565 
Consecutive  reactions,  289 
Conservation  of  mass,  18 
Constant,  equilibrium,  184 

ion-product,  470 

ionization,  238 

molecular  depression,  214 
Constituents,  8 
Contact  action,  definition,  29 
Copper,  501 

compounds,  504 

pyrites,  264 

refining,  511 
Copperas,  634 
Corn  syrup,  404 
Cordite,  359 

Corrosive  sublimate,  533 
Coulomb,  237 
Couples,  549 
Cream  of  tartar,  463 
Critical  temperature,  78 
Cryolite,  204,  554 
Crystal  structure,  306 
Crystallization,  water  of,  96 
Crystals,  82 
Cupellation,  513 
Cupric  bromide,  249 

chloride,  534 

nitrate,  211 

oxide,  34,  507 

sulphate,  95,  509 
Cuprous  chloride,  504 
Cyanogen,  420 

DAVY,  15 

Deacon's  process,  156,  177 
Decantation,  12 
Decomposition,  14 
Decrepitation,  449 
Deliquescence,  134 
Density,  gases,  73 
relative,  of  gases,  152 

and  salts  under  the  positive  radical. 


INDEX 


655 


Density,  solutions,  138 

vapor,  74 
Depilatory,  486 
Dextrin,  404 
Dextrose,  404 
Dewar  flask,  335 
Dialysis,  416 
Diamond,  376 
Diffusion,  57,  78 
Digestion,  422,  423 
Dihydrol,  138 

Dimorphous  substances,  266 
Displacement,  55 

ionic,  258 
Dissociation,  93 

cases  of,  116 

in  solution,  210 
Distillation,  fractional,  390 
Double  decomposition,  147 

ionic  formulation  of,  251 
Drummond  light,  58 
Dust  in  air,  328 
Dyes,  411,  419,  563 
Dynamite,  359 

EARTHENWARE,  561 
Earths,  rnetals  of  the,  553 

rare,  304,  553 
Efflorescence,  96 
Electric  energy,  units,  539 
Electric  furnace,  363,  377,  380 
Electric  waves,  wireless,  303 
Electrolysis,  55,  155 

explanation  of,  232 

products  of,  227 

quantities  of  electricity,  237 
Electrolytic  refining,  511,  549,  574 
Electromotive  chemistry,  539 

series,  260,  547,  650 
Electrons,  322,  609 

and  ions,  235 
Electrophoresis,  417 
Electroplating,  510,  516 
Electrotyping.  510,  516 
Element,  used  in  two  senses,  16 
Elements,  common,  17 

non-metallic,  94,  296 

metallic,  94,  296 
Emeralds,  524 
Emulsion,  121,  418 
Energy,  chemical  change  and,  167 

conservation  of,  170 

internal,  172 

source  of  world's,  388 
Enzymes,  405 

***  Acids  are  all  listed  under  "  acid 


Equations,  44 

concurrent  reactions,  317 

partial,  194 

thermochemical,  174 

writing,  50,  52,  97,  115,  276,  322- 

326,  360 

Equilibrium,  chemical,  177 
characteristics,  179 

constant,  184 

displacement  of,  90,  143 

ionic,  238,  247,  466 
displacement  of,  248 
saturated  solutions,  469 

liquid  and  vapor,  89 

saturated  solution,  130 

three  characteristics  of,  89 
Equivalent  weights,  65 
Esters,  413 
Ethyl  acetate,  413 
Ethylene,  392,  393,  400 
Explanation,  22 
Explosives,  358 

FATS,  414,  423 
Fehling's  solution,  404,  507 
Felspar,  3,  430 
Fermentation,  406 
Ferric  compounds,  634 

thiocyanate,  182 
Ferrosilicon,  425 
Ferrous  compounds,  633 

sulphide  and  acids,  272 
Ferrovanadium,  594 
Fertilizers,  387,  451,  488 

ammonium  sulphate,  340 

nitrogen,  339 
Filter,  Pasteur,  92 
Filtration,  12,  91 
Fire-damp,  391 
Fire  extinguishers,  379,  384 
Fixation  of  nitrogen,  352 
Flame,  396 

blast-lamp,  398 

Bunsen,  398,  400 

cone-separator,  401 

luminosity,  399 

structure,  399 
Flotation,  froth,  503 
Flour,  wheat,  3 
Fluor-spar,  475 
Fluorine,  204 
Fluorite,  475 
Flux,  438 
Foods,  421-423 

fuel  value  of,  95 

and  salts  under  the  positive  radical. 


656 


INDEX 


Formula,  reaction,  95 
Formulae,  44 

and  valence,  63 

from  data,  45 

making,  97 

graphic,  291,  358,  372 

molecular,  109 

Formulation,  of  chemical  equilibrium, 
183 

of  double  decomposition,  251 

of  neutralization,  254 

of  precipitation,  252 
Fractions  ionized,  data,  241 
Freezing-point,  definition,  86 

of  solutions,  134,  213 
Freezing  mixtures,  134 
Froth  flotation,  503 
Fructose,  404 
Fuels,  410 
Furnace,  electric,  363,  377,  380,  488 

G.M.V.,  101,  103 
Galena,  573 
Gallium,  553 
Galvanized  iron,  528 
Garnet,  82 
Gas,  blau,  396 

coal,  409 

-lighters,  580 

oil,  396 

perfect,  79 

producer,  385 

water,  386 

carburetted,  395 
Gases,  density,  73 

laws  of,  70,  76 

liquefiabilities,  278 

liquefaction,  78,  334 

measurement,  70 

mixed,  72 

solubilities  of,  278 
Gasoline,  391 
German  silver,  504 
Germanium,  567 
Glass,  493 

etching,  206,  494 

quartz,  428 

uranium,  606 

water,  428 

Glauber's  salt,  96,  464 
Glucinum,  523 
Glucose,  403 
Gluten,  3 
Glycerine,  413 
Gypsum,  485 

***  Acids  are  all  listed  under  "acid 


Gold,  518 

compounds,  520 
Gram-molecular  volume,  102 
Granite,  2,  430 
Grape-sugar,  404 
Graphic  formulae,  291 
Graphite,  376,  377  . 

Guano,  338,  348 
Guncotton,  350,  358 
Gunpowder,  449 
Gypsum,  84,  485 

HALOGEN  family,  192,  207 

chemical  relations,  319 
Hardness,  scale  of,  648 
Heat,  animal,  36 

of  neutralization,  255 

of  solution,  125 

of  vaporization,  86 

thermochemistry,  174 
Heavy-spar,  496 
Helium,  15,  336,  608 
Household  ammonia,  344 
Humidity,  328 
Hydrates,  95 
Hydrazine,  340,  345 
Hydrocarbons,  389 

cracking  of,  395 

unsaturated,  392 
Hydrogen,  49 

chemical  properties,  58 

commercial  sources,  56 

dissociation  of,  113 

history,  49 

-ion,  246 

nascent,  360 

physical  properties,  56 

preparation,  49,  51,  53,  55,  56 
Hydrogen  bromide,  196 
Hydrogen  chloride,  properties,  144 

composition  by  volume,  164 

preparation,  141 
Hydrogen  iodide,  201 

peroxide,  222 

sulphide,  267 

and  iodine,  202 
Hydrolysis,  197,  437 

of  salts,  271      - 
Hydrolyte,  475 
Hydrone,  50 
Hydroxide-ion,  246 
Hydroxylamine,  340 
Hypo,  464 
Hypochlorites,  308 
Hypochlorous  anhydride,  307 
and  salts  under  the  positive  radical. 


INDEX 


657 


ICE,  85 

heat  of  fusion,  86 
Indicators,  257 
Indium,  553 
Infusorial  earth,  428 
Ink,  printers'  and  India,  398 

writing,  634 

Internal  rearrangement,  148,  421 
Invar,  632 
lodic  anhydride,  319 
Iodine,  198 

chlorides  of,  208 

union  of  hydrogen  and,  203 
lodoform,  392 
lodothyrene,  199 
Ion-product  constant,  470 
Ionic  equilibrium,  466 
Ionic  substances,  245 

names  of,  236 
lonization,  226 

activity  and,  242 

constant,  238 

degree  of,  240,  241 

oxidation  and,  322 

questions  answered,  234 
lonogens,  classes,  245 

non-ionic  formation,  260 
Ions,  and  electrons,  235 

migration  of,  229,  231 
Iridium,  645 
Iron,  625 

carbonyls,  638 

cast,  627 

chemical  properties,  626,  632 

compounds,  633 

galvanized,  528,  550 

metallurgy,  626 

passive,  633 

Russia,  632 

wrought,  628 
Isomers,  421 
Isoprene,  392 

JAVEL,  eau  de,  448 

KAINITE,  451 
Kaolin,  430,  561 
Kerosene,  391 
Kindling  temperature,  35 
Kipp  apparatus,  54 
Krypton,  337 

LACTOSE,  404 
Lakes,  565 
Lampblack,  398 

***  Acids  are  all  listed  under  "acid 


Lard,  414 
Laughing  gas,  358 
Lavoisier,  6,  15,  26 
Law,  21 

Avogadro's,  77 

Boyle's,  71,  76 

Charles',  72,  76 

combining  weights,  42 

conservation  of  mass,  18 

Dalton's,  72 

definite  proportions,  17,  614 

Dulong  and  Petit's,  108 

Faraday's,  232 

Gay-Lussac's,  98 

Henry's,  128 

Le  Chatelier's,  190 

mass  action,  182 

multiple  proportions,  47 
Law  of,  chemical  change,  7 

component  substances,  3 

molecular  concentration,  182 

partition,  129 
Law,  periodic,  300 

van't  Hoff's,  188 
Laws  of  gases,  deviations  from,  78, 

79 

Le  Blanc  process,  460 
Lead,  573 

compounds,  574 

from  radium,  613 

from  thorium,  614 

pencils,  378 

red,  575 

white,  577    - 
Lime,  477 

light,  58 
Liquids,  associated,  206 

molecular  relations,  81 
Litharge,  575 
Lithium,  465 
Litmus,  258 
Lithopone,  497 
Lomonssov,  4,  5,  15,  80 

MAGNALIUM,  555 
Magnesium,  524 

compounds,  525 

nitride,  339 
Malachite,  502 
Maltose,  404 
Manganese,  617 
Manganic  compounds,  620 
Manganin,  618 
Manganites,  621 
Manganous  compounds,  619 
and  salts  under  the  positive  radical. 


658 


INDEX 


Marsh  gas,  391 
Marsh's  test,  584 
Matches,  365 
Matrix,  264 

Matter,  structure  of,  74 
Maypw,  5,  14,  25 
Melting-point,  definition,  86 
Mendelejeff,  298 
Mercuric  oxide,  14,  27 
Mercury,  532 

compounds,  533 
Metallic  elements,  94 

chemical  relations,  436 
Metals,  electromotive  series  of,  260 

extraction,  438 

melting-points,  435 

occurrence,  437 

order  of  activity,  59,  650 

physical  properties,  434 

potential  differences,  547 

world's  production,  436 
Methane,  378,  391 
Methyl  orange,  258 
Methylated  spirit,  407 
Metric  system,  648 
Mica,  2,  430 
Microcpsmic  salt,  371 
Migration  of  ions,  229 
Mill,  422 
Minium,  575 
Mirrors,  silvering,  517 
Mixture,  4 
Moisture,  surface,  88 
Molar  solutions,  125 
Molar  weight,  102 
Molasses,  405 
Mole,  102 

number  of  molecules  in,  103 
Molecular  equations,  interpretations, 

115 

Molecular  formulae,  109 
Molecular  theory,  74 

gases,  74 

histpry,  80 

liquids,  81 

liquid  and  vapor,  88 

of  solutions,  125 

solids,  81 
Molecular  weights,  100 

by  freezing-point,  134 

in  solution,  214 

of  elements,  110 
Molybdenum,  604 
Monde  process,  642 
Monel  metal,  641 


Mortar,  478 

Moseley's  atomic  numbers,  303 

NAPHTHA,  391 

Naphthalene,  411 

Neon,  15,  337 

Neutralization,  formulation  of,  254 

heat  of,  255 
Nickel,  640 

carbonyl,  642 

sulphate,  83 
Niton,  337,  612 
Nitric  anhydride,  349 
Nitric  oxide,  350 
Nitrides,  339 
Nitro-lime,  487 
Nitrogen,  338 

iodide,  346 

tetroxide,  351 

trichloride,  346 
Nitroglycerine,  350,  358 
Nitrosyl  chloride,  356 
Nitrous  anhydride,  357 

oxide,  357 

Nomenclature,  64,  306 
Non-metallic  elements,  94 

potential  differences,  548 
Normal  solutions,  124 

OIL  gas,  396 

Oil,  cotton  seed,  414 

of  vitriol,  285 

olive,  414 
Oleum,  285 

Open-hearth  process,  630 
Osmium,  644 

Osmotic  pressure,  125,  135 
Ostwald,  21 
Oxidation,  36 

always  with  reduction,  269 

and  reduction,  320-326 
Oxides,  acidic,  94 

basic,  94 

order  of  stability,  60 
Oxidizing  agents,  explanation  of  ac- 
tivity, 221 
Oxygen,  25 

chemical  properties,  31 

history  of,  25 

physical  properties,  30 

preparation  of,  26 

uses  of,  37 

why  atomic  weight  16,  46 
Oxone,  28 
Ozone,  219 


***  Acids  are  all  listed  under  "acid"  and  salts  under  the  positive  radical. 


INDEX 


659 


PAINT,  579 

Hthopone,  497 

luminous,  486 

permanent  white,  496 
Palladium,  57,  645 
Paper,  402 

sizing,  560 
Paraffin,  390 
Paris  green,  508 
Parke's  process,  513 
Pasteur  filter,  92 
Pauling  process,  353 
Pearl  ash,  450 
Perchlorates,  315 
Perchloric  anhydride,  316 
Periodic  system,  297,  inside  rear  cover 
Permutite,  491 
Peroxidates,  223 
Peroxides,  223 
Petrol,  391 
Petrolatum,  391 
Petroleum,  390 

refining  of,  393 
Phenolphthalein,  258 
Phosgene,  163 
Phosphate  rock,  362 
Phosphine,  365 
Phosphonium  iodide,  366 
Phosphorescence,  364 
Phosphoric  anhydride,  367 
Phosphorite,  362,  486 
Phosphorus,  362 

acids  of,  368 

pentachloride,  117,  367 

pentasulphide,  373 

pentoxide,  367 

trichloride,  367 

tribromide,  197 

tri-iodide,  201 

vapor,  117 

Photography,  517,  600,  637 
Picture  restoring,  224 
Plants  and  carbon  dioxide,  387 
Plaster  of  Paris,  485 
Plastics,  359 
Platinum,  646 

as  catalyst,  59 
Plumbago,  375 
Polarization,  548 
Polonium,  614 
Polysulphides,  274 
Porcelain,  561 
Potash,  450 
Potassium,  443 

alum,  558 

***  Acids  are  all  listed  under  "  acid ' 


Potassium,  bisulphite,  452 

bitartrate,  463 
Potassium  bromide,  445 

sulphuric  acid  on,  196 
Potassium  carbonate,  450 

chlorate,  27,  313,  448,  469 

chloride,  444 

chromate,  597 

cobalticyanide,  640 

cobaltinitrite,  640 

cuprocyanide,  508 

cyanate,  421 

cyanide,  451 

cGchromate,  597 

ferricyanide,  637 

ferrocyanide,  451,  636 

fluorides,  445 

hydroxide,  446 

hypochlorite,  308 
Potassium  iodide,  445 

sulphuric  acid  on,  201 
Potassium  manganate,  621 

nitrate,  83,  448 

oxides,  447 

perchlorate,  315,  448 

permanganate,  83,  157,  225,  621 

sulphate,  451 

sulphides,  452 

thiocyanate,  421 

tri-iodide,  200 

Potential  differences,  single,  546,  547 
Potentials,  discharging,  548 
Precipitates,  description  of,  147 
Precipitation,  in  presence  of  acids,  484 

formulation  of,  252 

theory  of,  478 
Pressure,  osmotic,  126,  135 

partial,  72 

solution,  128 

vapor,  87 
Priestley,  14,  26 
Problems,  arithmetical,  45,  66,  115, 

149 

Producer  gas,  385 
Properties,  specific  chemical,  30 

specific  physical,  19,  30,  31 
Proteins,  422,  486 

test,  350 

Prussian  blue,  637 
Pyrene,  379 
Pyrite,  264,  275,  636 

QUALITATIVE  analysis,  537 
Quantitative  experiments,  33,  34,  35 
Quartz,  3,  83,  427 
and  salts  under  the  positive  radical. 


660 


INDEX 


Quartz  glass,  428 
Quicklime,  477 

RADICALS,  53,  212 

positive  and  negative,  53 

valence  of,  62 
Radioactive  elements,  606 
Radioactivity,  significance,  614 

uranium  group,  612 
Radium,  608 
Reactions,  concurrent,  315,  317 

consecutive,  289 
Reaction  formula,  95 
Realgar,  586 

Reduction  and  oxidation,  320-326 
Refrigeration,  342 
Relations,  chemical,  163,  192 
Reversible  actions,  93 
Rey,  Jean,  5 
Rhodium,  645 
Rittman's  process,  391,  395 
Roasting,  275 
Rochelle  salts,  463 
Rock  crystal,  see  Quartz 
Root  nodules,  339 
Rubber,  synthetic,  392 
Rubidium,  453 
Rusting,  1,  4,  5,  6 
Ruthenium,  644 

SALAMMONIAC,  453 
Saleratus,  450 
Salt,  common,  82 
Saltpeter,  air,  353 

Bengal,  347 

Chile,  347,  448 
Salts,  149 

double,  245 

fractions  ionized,  242 

ions  of,  246 

non-ionic  formation,  261 

mixed,  245 

properties  in  solution,  210 
Sandstone,  430 
Saponification,  415 
Saponin,  420 
Scheele,  26 
Schoenite,  451 
Schwerin  process,  562 
Selenite,  485 
Selenium,  293 
Sewage,  36 

Siemens-Martin  process,  630 
Silicates,  430 
Silicon,  425 


Silicon,  dioxide,  427 

tetrafluoride,  206,  427 
Silk,  imitation,  359,  507 
Silver,  512 

complex  compounds,  514 

salts,  515 
Slag,  438 

Smokeless  powder,  359 
Soap,  412,  415,  490 

cleansing  power,  418 

salting  out,  417 
Soda,  washing,  96 
Soda-water,  382 
Sodium,  457 

-ammonium  phosphate,  371 

bicarbonate,  462 

carbonate,  96,  460 

chloride,  82,  458,  472 

cyanide,  488 

cGchromate,  598 

hydride,  458 

hydroxide,  459 

hyposulphite^  464 

iodate,  318 

metaphosphate,  371 

nitrate,  459 

orthophosphates,  370,  464 

oxides,  459 

palmitate,  412 

peroxide,  28,  222 

persulphate,  291 

silicate,  428 

sulphate,  96,  463 
solubilities,  132 

tetraborate,  432,  465 
Sodium  thiosulphate,  290,  464 
Solids,  molecular  relations,  81 
Solubilities,  131,  inside  front  cover 
Solubility,  gases,  128 

measurement  of,  123 

product,  471 

temperature  and,  130 

units  to  express,  124 
Solution,  121 

as  a  process,  127 

dissociation  in,  210 

freezing-points,  134,  213 

heat  of,  125 

insoluble  salts  by  acids,  481 

molecular  theory,  125 

physical  or  chemical,  138 

pressure,  128 

rule  for,  479 

saturated,  123,  128 
definition,  133 


***  Acids  are  all  listed  under  "acid"  and  salts  under  the  positive  radical. 


INDEX 


661 


Solution,  solid,  122 

supersaturated,  132 

volume  changes  in,  138 
Solutions,  boiling-points,  135 

see  Colloids 

densities,  138 

freezing-points,  134 

molar,  125 

normal,  124 

standard,  257,  623 

vapor  tension,  134 
Solvay  process,  461 
Specific  heat,  85,  108 
Spintharoscope,  608 
Sponges,  428 
Stability,  chemical,  38 

compounds,  93 
Stalactites,  476 
Starch,  3,  403,  423 
States  of  matter,  86 
Stationary  layers,  328,  398 
Steam,  86 
Stearin,  414 
Steel,  629-632 

alloys,  632 

chromium-vanadium,  632 

manganese,  618,  632 

nickel,  632 

tungsten,  606 
Stereotype  metal,  511 
Stibine,  588 
Strontium,  495 
Structure  of  matter,  74 
Sublimation,  199 
Substance,  2 

simple,  15 
Substitution,  162 
Sucrase,  406 
Sucrose,  404 
Sugar,  84 

cane,  404 

invert,  405 
Sugars,  404 
Sulphates,  288 
Sulphate-ion,  287 
Sulphides,  270 

action  of  acids  on,  271 

insoluble,  classification  of,  273 

solubilities  of,  531 
Sulphites,  290 
Sulphur,  264 

properties,  266 

vapor,  117 
Sulphur,  acids  of,  280 

dioxide,  35,  275,  277 

***  Acids  are  all  listed  under  "acid 


Sulphur,  monochloride,  291 

trioxide,  279 

Sulphuryl  chloride,  278,  291 
Superphosphate,  487 
Symbols,  44 

T.  N.T.,  349 
Tanning,  chrome,  600 
Tantalum,  594 
Tartar-emetic,  589 
Tellurium,  294 

Temperature,  and  speed  of  reaction, 
59,  187 

conversion  table,  649 

critical,  78 
Tempering,  631 
Tensile  strength,  629 
Tension,  aqueous,  87,  649 
Thallium,  553 
Theory,  molecular,  74,  80 
Thermite,  556 
Thermochemistry,  174 
Thorium,  580 

oxide,  397 
Tin,  567 

compounds,  570 

-plate,  550,  569 

see  Stannous  and  Stannic 
Titanium,  580,  630 
Titration,  256 
Toluene,  349,  392 
Transition  points,  86 
Trinitrotoluene,  349 
Tungsten,  605 
Turnbull's  blue,  637 

ULTRAMARINE,  563 
Ultramicroscope,  416 
Units,  electrical,  237 

of  measurement,  648 
Uranium,  606 

radioactivity  of,  612 
Urea,  421 

VALENCE,  61 

and  formulae,  63 

and  oxidation,  321 

definition,  62 

exceptional,  64 

how  ascertained,  63 
Vanadium,  593 
Vapor,  density,  measurement  of,  74 

equilibrium  with  liquid,  89 

pressure,  87 

saturated,  88 
and  salts  under  the  positive  radical. 


662 


INDEX 


Varnish,  black,  398 

Vaseline,  391 

Ventilation,  328 

Verdigris,  508 

Vermilion,  535 

Vitriols,  529 

Volt,  237 

Volume,  gram-molecular,  102 

Volumetric  analysis,  257 

WARNINGS,  68,   102,   111,   119,   133, 

175,  260 

Washing  soda,  96,  462 
Water,  85 

as  solvent,  90 

chemical  properties,  92 

coagulation  process,  560 

composition  of,  98 

dihydrol,  138 

domestic,  purification,  312 
Water  gas,  386 

carburetted,  395 

***  Acids  are  all  listed  under  "  acid 


Water  glass,  428 
hard,  415,  489-493 
of  crystallization,  96 
physical  properties,  85 
vapor  tensions,  649 

Waters,  natural,  91 

Weights,  atomic,  103 
equivalent,  65 
molar,  102 
molecular,  100 

Welsbach  lamp,  397 
mantles,  580 

Whisky,  406 

Witherite,  496 

Wood,  distillation  of,  408 

Wood's  metal,  591 

X-RAYS,  303,  609 
Xenon,  337 

ZINC,  527 

compounds,  528 
Zymase,  406 
and  salts  under  the  positive  radical. 


i 


UNIVERSITY  OF  CALIFORNIA  LIBRARY 

This  book  is  DUE  on  the  last  date  stamped  below. 


OCT  14  1947 


REC'D  LD 
OCT  13  1356 


MAR  15 


1{I?V2319630S! 


2lOct'53Vn 
OCT1  9  19 


LD  21-100m-12,'46(A2012sl6)4120 


THE  UNIVERSITY  OF  CALIFORNIA  LIBRARY 


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